BIOLOGY 
t\BRARY 


OUTLINES   OF   CHEMISTRY 


THE  MACMILLAN  COMPANY 

NEW  YORK    •    BOSTON   .    CHICAGO 
SAN   FRANCISCO 

MACMILLAN  &  CO.,  LIMITED 

LONDON   •    BOMBAY    •    CALCUTTA 
MELBOURNE 

THE  MACMILLAN  CO.  OF  CANADA,  Lm 

TORONTO 


OUTLINES   OF  CHEMISTRY 


A  TEXT-BOOK  FOR  COLLEGE  STUDENTS 


BY 


LOUIS    KAHLENBERG,   PH.D. 

PROFESSOR    OF    CHEMISTRY    AND    DIRECTOR    OF    THE    COURSE 
IN    CHEMISTRY    IN    THE    UNIVERSITY   OF    WISCONSIN 


gorfc 

THE   MACMILLAN  COMPANY 
1912 

All  rights  reserved 


COPYRIGHT,  1909, 
BY  THE  MACMILLAN  COMPANY. 


Set  up  and  electrotyped.     Published  September,  1909.     Reprinted 
January,  September,  October,  1910. 

Reprinted  with  corrections,  January,  September,  1911 ;  August,  1912. 
October,  1912. 


Nmrfaoob 

J.  8.  Gushing  Co.  —  Berwick  &  Smith  Co. 
Norwood,  Mass.,  U.S.A. 


PREFACE 

• 

THIS  book  is  intended  to  represent  one  year's  work  of 
chemistry  in  college.  It  should  be  used  in  connection  with 
a  course  of  experimental  lectures  and  laboratory  exercises. 
The  matter  has  been  selected  so  as  to  meet  the  needs  of  those 
that  can  devote  but  one  year  to  the  study  of  chemistry,  and 
also  to  serve  as  a  suitable  basis  for  future  work  in  the  case  of 
students  who  desire  to  pursue  the  subject  further.  In  writing 
the  book,  the  author  has  naturally  had  in  mind  the  needs  of 
his  own  students,  over  six  hundred  in  number,  who  are  pre- 
paring for  careers  in  chemistry,  pharmacy,  medicine,  engineer- 
ing, or  agriculture,  or  who  desire  a  course  in  chemistry  for  work 
in  other  natural  sciences  or  as  a  means  of  general  culture. 

In  the  first  five  chapters,  experimental  work  has  been  placed  in 
the  foreground,  and  all  reference  to  atomic  and  molecular  theo- 
ries has  been  purposely  avoided  in  order  that  the  student  may 
properly  be  impressed  with  the  fundamental  facts  and  laws, 
which  are  independent  of  the  theories,  though  they  serve  as 
a  foundation  for  the  latter.  In  the  sixth  chapter,  these  funda- 
mental laws  are  then  reviewed,  and  the  atomic  and  molecular 
theories  are  presented  as  views  growing  out  of  the  experimental 
facts.  The  nomenclature  is  then  also  introduced,  and  the 
reactions  which  so  far  have  been  written  in  words  are  expressed 
by  means  of  chemical  symbols.  This  offers  an  excellent  oppor- 
tunity for  reviewing  the  experimental  work  of  the  foregoing 
chapters.  While  the  teacher  is  somewhat  inconvenienced  by 
thus  postponing  the  introduction  of  the  atomic  theory  and  the 
use  of  formulation  till  the  student  has  at  least  a  fair  stock 
of  carefully  selected  facts  upon  which  to  found  the  theory, 
it  really  pays  to  make  the  exertion,  for  thus  greater  interest 
is  created  and  the  student  sees  the  facts  and  theoretical  views 

T 

266991 


yi  PREFACE 

in  their  proper  relations.  He  becomes  a  clear,  logical  thinker, 
and  does  not  look  upon  the  atomic  and  molecular  theories  as 
something  arbitrary,  metaphysical,  and  well-nigh  incomprehen- 
sible. The  method  here  adopted  is  not  new.  It  is  essentially 
the  same  in  principle  as  that  followed  by  Bunsen  and  many 
other  successful  teachers  of  chemistry. 

Throughout  the  book,  the  endeavor  has  been  to  convey  the 
salient  facts  in  as  simple  and  direct  a  manner  as  possible, 
developing  cardinal  principles,  and  carefully  keeping  the  dis- 
tinction between  facts  and  theories  in  mind.  The  aim  has 
been  to  enlist  the  interest  of  the  student  in  the  study  of  chem- 
istry, and  to  this  end  the  historical  development  of  certain 
aspects  of  the  subject  has  been  presented  as  far  as  space  would 
permit. 

The  most  important  technical  applications  and  processes 
have  constantly  been  emphasized,  though  they  have  been  in- 
troduced in  connection  with  the  description  of  the  various 
elements  and  compounds  rather  than  as  special  chapters.  On 
the  other  hand,  it  has  been  thought  best  to  treat  the  subjects 
of  thermochemistry  and  solutions  and  electrolysis  in  special 
chapters,  after  a  sufficient  number  of  fundamental  facts  have 
been  acquired  by  the  student,  so  that  he  is  in  a  position  to 
comprehend  the  more  difficult  relationships  which  these  topics 
involve. 

Only  the  essential  parts  of  chemical  theory  which  can  be 
comprehended  by  college  students  who  are  beginning  the  study 
of  chemistry  have  been  presented.  My  own  experience  would 
indicate  that  fully  as  much  has  been  given  as  they  can  well 
digest  at  this  stage  of  their  work.  In  touching  upon  contro- 
verted points,  the  aim  has  been  to  present  both  sides  of  the 
question  involved.  I  have  felt  that  the  teacher  should  not 
entirely  avoid  mooted  questions  even  during  the  first  year  of 
work  in  chemistry,  for  by  so  doing  the  impression  is  conveyed 
that  all  matters  are  in  a  settled  state,  and  thus  a  powerful 
stimulus  toward  further  study  and  inquiry  is  lost.  On  the 
whole,  however,  the  presentation  of  the  subject  has  been  along 
rather  well  established,  conservative  lines.  The  dominant  idea 
has  been  to  select$with  care  what  the  student  needs,  what  he 
can  reasonably  be  asked  to  comprehend,  at  his  stage  of  advance- 
ment, and  to  present  this  in  a  clear,  simple,  and  direct  manner, 


PREFACE  vil 

taking  the  trouble  to  repeat  and  to  emphasize  here  and  there 
in  order  to  secure  the  desired  end. 

My  best  thanks  are  due  to  Dr.  J.  H.  Walton  for  suggestions 
and  reading  of  proof,  also  to  Messrs.  C.  W.  Hill,  D.  Klein, 
F.  C.  Krauskopf,  and  W.  G.  Wilcox  for  reading  proof  sheets 
of  some  of  the  chapters.  Additional  suggestions  or  corrections 
to  be  used  in  preparing  further  editions  will  be  welcomed  from 

others. 

LOUIS  KAHLENBERG. 
MADISON,  WISCONSIN, 
June  8, 1909. 


CONTENTS 


CHAPTER  I 

THE  SCOPE  OF  CHEMISTRY  AND  ITS  RELATIONS  TO  OTHER  SCIENCES  — 
CHEMICAL  CHANGE,  ELEMENTS,  AND  COMPOUNDS 

PAGE 

Physical  and  Chemical  Changes  —  Definite  Proportions  —  Solutions 
and  Chemical  Compounds  —  Chemical  Elements  —  Compounds  — 
Types  of  Chemical  Change  —  Conservation  of  Mass — Conserva- 
tion of  Energy  —  Cause  of  Chemical  Change  —  Chemical  Affinity 

—  Factors  affecting  Chemical  Change 1 

CHAPTER  II 
HYDROGEN 

History  —  Occurrence  —  Preparation  —  Properties  —  Uses — Hydrogen 

Equivalents  of  the  Metals 13 

CHAPTER  III 
OXYGEN 

History  —  Occurrence  —  Preparation  —  Properties  —  Combustion  in 
the  Air  —  Kindling  Temperature  and  Temperature  of  Combustion 

—  Heat  of  Combustion  —  Different  Stages  of  Oxidation  —  Law  of 
Multiple   Proportions  —  Role  of  Oxygen  in    Respiration  —  Oxy- 
hydrogen  Blowpipe  —  Detonating  Gas  —  Combustion  of  Oxygen 

in  Hydrogen  —  Earlier  Views  of  Combustion 24 

CHAPTER  IV 
WATER 

Occurrence  —  Preparation — Natural  Waters  —  Potable  Water  —  Min- 
eral Water —  Composition  —  Gay-Lussac's  Law  of  Combination  of 
Gases  by  Volume  —  Properties  of  Water  —  Super-cooled  Water  — 
Change  of  Freezing-point  with  Pressure  —  Principle  of  Le  Cha- 
telier  —  Crystalline  Nature  of  Ice  —  Compounds  with  Water  — 

Water  as  a  Solvent 36 

ix 


X  CONTENTS 

CHAPTER  V 
HYDROCHLORIC  ACID  AND  CHLORINE 

PAOB 

Preparation  and  Properties  of  Hydrochloric  Acid  —  Composition  and 
Chemical  Behavior  of  Hydrochloric  Acid  —  Occurrence,  History, 
and  Properties  of  Chlorine  —  Uses  of  Chlorine  —  Some  Compounds 
of  Chlorine  with  Oxygen  —  Law  of  Reciprocal  Proportions  .  .  49 

V 

CHAPTER  VI 

THE  LAWS  OF  COMBINING  WEIGHTS  AND  COMBINING  VOLCTMES 
AND  THE  ATOMIC  AND  MOLECULAR  THEORIES 

Retrospect  —  Laws  of  Definite,  Multiple,  and  Reciprocal  Proportions  — 
Combining  Weights  and  Chemical  Equivalents  —  Chemical  Sym- 
bols—  Atomic  Theory  of  Matter  —  Difference  between  Theory 
and  Law  —  Law  of  Combination  of  Gases  by  Volume  —  Avogadro's 
Hypothesis — Molecular  Weight  Determinations  —  Determination 
of  Atomic  Weights  —  Law  of  Dulong  and  Petit  —  Other  Methods 
of  Choosing  Atomic  Weights  from  the  Combining  Weights  — 
Law  of  Isomorphism  —  Table  of  Atomic  Weights  —  Interpreta- 
tion of  a  Chemical  Formula  —  Valence  and  Structural  Formulae 

—  Nomenclature  —  Chemical    Equations  —  Retrospect  —  Phe- 
nomena of  the  Nascent  State 58 

CHAPTER   VII 
OZONE,  ALLOTROPY,  AND  HYDROGEN  PEROXIDE 

History,  Occurrence,  and  Preparation  of  Ozone — Relation  between 
Ozone  and  Oxygen  —  Allotropy  —  Properties  of  Ozone — History, 
Occurrence,  and  'Preparation  of  Hydrogen  Peroxide  —  Properties 
of  Hydrogen  Peroxide  —  Formula  of  Hydrogen  Peroxide  —  Uses 
of  Hydrogen  Peroxide  —  Ozonic  Acid 88 

CHAPTER  VIII 
THE  HALOGENS 

The  Halogen  Family  —  Compounds  of  Chlorine  with  Oxygen  —  Hypo- 
chlorous  Acid  and  Hypochlorites  —  Chloric  Acid  and  Chlorates  — 
Perchloric  Acid  and  Perchlorates  —  Nomenclature  and  General 
Relations  —  Occurrence,  Preparation,  and  Properties  of  Fluorine 
— 'Hydrofluoric  Acid  —  Occurrence,  Preparation,  and  Properties 
of  Bromine  —  Hydrobromic  Acid  —  Oxy-acids  of  Bromine  — 
Bromic  Acid  and  Bromates  —  Uses  of  Bromine  and  its  Compounds 

—  History  and  Occurrence  of  Iodine  —  Preparation  of  Iodine  — 


CONTENTS  xi 

PAGE 

Properties  of  Iodine  —  Uses  of  Iodine  —  Hydriodic  Acid  —  Oxide 
of  Iodine  —  Oxy-acids  of  Iodine  —  Compounds  of  the  Halogens 
with  Each  Other  —  General  Relations  of  the  Halogens  to  One 
Another 96 

CHAPTER  IX 

ACIDS,  BASES,  SALTS,  HYDROLYSIS,  MASS  ACTION,  AND 
CHEMICAL  EQUILIBRIUM 

Acids  —  Bases  —  Salts  —  Older  View  of  the  Process  of  Salt  Formation 
—  Acid-  and  Base-forming  Elements  —  Other  Views  of  Solutions 
of  Acids,  Bases,  and  Salts  —  Basicity  of  Acids  —  Acid  Salts  — 
Acidity  of  Bases  —  Basic  Salts  —  Normal  Salts  —  Acidimetry  and 
Alkalimetry  —  Indicators  —  Hydrolysis  —  Mass  Action  —  Chemi- 
cal Equilibrium  —  Additional  Illustrations  of  Chemical  Equi- 
librium and  the  Operation  of  the  Law  of  Mass  .Action  —  Strength 
of  Acids  and  Bases 120 

CHAPTER  X 

NITROGEN,  THE  ATMOSPHERE,  AND  THE  ELEMENTS  OF  THE  HELIUM 

GROUP 

History  and  Occurrence  of  Nitrogen  —  Preparation  and  Properties  of 

Nitrogen  —  The  Air  —  The  Elements  of  the  Helium  Group  .        .     139 

CHAPTER  XI 

COMPOUNDS  OF  NITROGEN  WITH  HYDROGEN  AND  WITH  THE 
HALOGENS 

History  and  Occurrence  of  Ammonia  —  Preparation  and  Properties  of 
Ammonia  —  Hydrazine  —  Hydroxylamine  —  Hydrazoic  Acid  — 
Compounds  of  Nitrogen  with  the  Halogens 150 

CHAPTER  XII 
OXY-ACIDS  AND  OXIDES  OF  NITROGEN 

History,  Occurrence,  and  Preparation  of  Nitric  Acid  —  Properties  of 
Nitric  Acid  —  Nitrogen  Pentoxide  —  Nitric  Oxide  —  Nitrogen 
Dioxide  and  Tetroxide  —  Nitrous  Acid  —  Nitrogen  Trioxide  — 
Hyponitrous  Acid  —  Nitrous  Oxide  —  General  Considerations  .  161 

CHAPTER  XIII 
SULPHUR,  SELENIUM,  AND  TELLURIUM 

Occurrence  and  Preparation  of  Sulphur  —  Properties  of  Sulphur  — 
Uses  of  Sulphur  —  Crystals  and  Crystal  Systems  —  Hydrogen 


xii  CONTENTS 


Sulphide  —  Poly-sulphides  and  Hydrogen  Persulphide  —  Compari- 
son of  Hydrogen  Sulphide  with  Water  —  Compounds  of  Sulphur 
with  the  Halogens — Sulphur  Dioxide  and  Sulphurous  Acid  — 
Sulphur  Sesquioxide  —  Sulphur  Trioxide  and  the  Contact  Process 
of  making  Sulphuric  Acid  —  Sulphuric  Acid  and  the  Lead  Cham- 
ber Process  —  Properties  of  Sulphuric  Acid  —  Hydrates  of  Sul- 
phuric Acid —  Pyrosulphuric  Acid — Thiosulphates — Persulphates 
—  Polythionic  Acids  —  Thionyl  Chloride  —  Sulphuryl  Chloride, 
Selenium  —  Compounds  of  Selenium  —  Tellurium  —  Compounds 
of  Tellurium  —  General  Considerations 176 


CHAPTER  XIV 
CARBON  AND  SOME  OF  ITS  TYPICAL  COMPOUNDS 

Occurrence  and  Allotropic  Forms  of  Carbon  —  Chemical  Behavior  of 
Carbon  —  Carbon  Dioxide  —  Properties  of  Carbon  Dioxide  — 
Physiological  Effects  of  Carbon  Dioxide  —  Relations  of  Carbon 
Dioxide  to  Plant  and  Animal  Life  —  Early  Work  on  Carbon 
Dioxide  —  Carbon  Monoxide  —  Properties  of  Carbon  Monoxide  — 
Physiological  Effects  of  Carbon  Monoxide  —  Carbon  Bisulphide 
—  Cyanogen  —  Hydrocyanic  Acid  —  Cyanates  and  Sulphocyauates  210 

CHAPTER  XV 
HYDROCARBONS  AND  ADDITIONAL  COMPOUNDS  OP  CARBON 

Hydrocarbons  —  General  Behavior  of  Hydrocarbons  —  Halogen  Sub- 
stitution Products  —  Alcohols  —  Phenols  —  Aldehydes  —  Organic 
Acids  —  Esters  —  Ethers  —  Ketones  —  Carbohydrates  —  Fermen- 
tation and  Enzymes  —  Starch  and  Dextrine  —  Cellulose  —  Nitro- 
benzene, Aniline,  and  Coal  Tar  Dyes  —  Alkaloids  —  Proteins  .  232 

CHAPTER  XVI 

ILLUMINATING  GAS  AND  FLAMES 

Illuminating  Gas  —  Flame  —  Luminosity  of  Flame  —  Structure   of 

Flame  —  Davy  Safety  Lamp 265 

CHAPTER  XVII 
THERMOCHEMISTRY 

General  Remarks  —  Calorimeters  —  Laws  of  Thermochemistry  — 
Thermochemical  Equations  —  Thermochemical  Data  —  Tables  — 
Uses  of  Thermochemical  Data  275 


CONTENTS  xiii 

CHAPTER  XVIH 
SILICON  AND  BORON  AND  THEIR  IMPORTANT  COMPOUNDS 

PAGE 

Occurrence,  Preparation,  and  Properties  of  Silicon — Silicon  Dioxide 

—  Silicic  Acids — Action  of  Water  on  Silicates  —  Decomposition 
of  Silicates  in  the  Laboratory  —  Hydrogen  Silicide  —  Compounds 
of  Silicon  with  the  Halogens  —  Esters  of  Silicic  Acid  —  Silicon 
Carbide  —  Titanium  —  Zirconium  —  Thorium  —  Occurrence, 
Preparation,  and  Properties  of  Boron  —  Boric  Acid  and  its  Salts 

—  Other  Compounds  of  Boron 290 

CHAPTER  XIX 
PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH 

Occurrence  and  Preparation  of  Phosphorus  —  Properties  and  Allo- 
tropic  Forms  of  Phosphorus  —  Uses  of  Phosphorus,  Matches  — 
Compounds  of  Phosphorus  with  the  Halogens  —  Oxides  and  Acids 
of  Phosphorus  —  Formulae  of  the  Acids  of  Phosphorus  —  Com- 
pounds of  Phosphorus  with  Sulphur  —  Occurrence,  Preparation, 
and  Properties  of  Arsenic  —  Arsine  —  Compounds  of  Arsenic  with 
the  Halogens  —  Oxides  and  Oxy-acids  of  Arsenic  —  Occurrence, 
Preparation,  and  Properties  of  Antimony  —  Stibine  —  Compounds 
of  Antimony  and  Sulphur  —  Occurrence,  Preparation,  and  Prop- 
erties of  Bismuth  —  Halogen  Compounds  of  Bismuth  —  Oxides  of 
Bismuth  —  Bismuth  Salts  of  Oxy-acids  —  Bismuth  Trisulphide  — 
General  Considerations  of  the  Group  —  Vanadium,  Columbium, 
and  Tantalum 304 

CHAPTER  XX 
CLASSIFICATION  OF  THE  ELEMENTS  —  THE  PERIODIC  SYSTEM  .        .    337 

CHAPTER  XXI 
THE  ALKALI  METALS 

Occurrence,  Preparation,  and  Properties  of  Potassium  —  Potassium 
Hydride  —  Compounds  of  Potassium  with  the  Halogens  —  Potas- 
sium Hydroxide  —  Potassium  Oxide  —  Potassium  Chlorate  —  Po- 
tassium Nitrate — Potassium  Cyanide  —  Potassium  Carbonate  — 
Potassium  Silicate  —  Potassium  Fluosilicate  —  Potassium  Phos- 
phates—  Potassium  Sulphate  —  Potassium  Sulphite  —  Sulphides 
of  Potassium  —  Tests  for  Potassium  —  Rubidium  and  Caesium  — 
Occurrence,  Preparation,  and  Properties  of  Sodium  —  Sodium 
Chloride  —  Oxides  and  Hydroxides  of  Sodium  —  Sodium  Carbon- 
ate —  Sodium  Nitrate  —  Phosphates  of  Sodium  —  Sodium  Sul- 


xiv  CONTENTS 


phate  —  Sodium  Sulphite  —  Sodium  Thiosulphate  —  Sodium  Sili- 
cate —  Sodium  Cyanide  —  Sodium  Borate  —  Lithium  and  its 
Compounds  —  The  Alkali  Metals  as  a  Group  —  Spectrum  Analysis 

—  Ammonium  Salts  —  Detection  of  Ammonium  Salts    .        .        .    343 

CHAPTER  XXII 
THE  ALKALINE  EARTH  METALS 

Occurrence,  Preparation,  and  Properties  of  Calcium  —  Calcium  Oxide 

—  Cement  —  Calcium   Sulphate  —  Calcium   Sulphite  —  Calcium 
Sulphide  —  Calcium   Fluoride  —  Calcium   Chloride  —  Bleaching 
Powder  —  Calcium    Phosphate  —  Calcium    Carbide  —  Calcium 
Phosphide  —  Calcium   Cyanamide  —  Calcium    Silicide  —  Calcium 
Silicate  —  Glass  —  Occurrence,   Preparation,   and    Properties   of 
Strontium  —  Strontium    Compounds  —  Occurrence,   Preparation, 
and  Properties  of  Barium  —  Compounds  of  Barium  —  Detection 

of  the  Alkaline  Earth  Metals  —  Radium  and  Radio-activity  .        .    374 

CHAPTER  XXIII 
THE  METALS  OF  THE  MAGNESIUM  GROUP 

Glucinum  —  Occurrence,  Preparation,  and  Properties  of  Magnesium  — 
Magnesium  Oxide  —  Magnesium  Carbonate — Magnesium  Chloride 

—  Magnesium  Sulphate —  Magnesium  Phosphates  —  Magnesium 
Ammonium  Arsenate — Tests  for  Magnesium  —  Occurrence,  Prepa- 
ration, and  Properties  of  Zinc  —  Zinc  Oxide  —  Zinc  Carbonate  — 
Zinc    Chloride  —  Zinc    Sulphate  —  Zinc   Sulphide  —  Analytical 
Tests  for  Zinc  Salts  —  Occurrence,  Preparation,  and  Properties  of 
Cadmium  —  Cadmium  Compounds  —  Occurrence,  Preparation,  and 
Properties  of  Mercury  —  Amalgams — Compounds  of  Mercury  — 
Oxides  of  Mercury  —  Halides  of  Mercury  —  Mercuric  Cyanide  — 
Nitrates  of  Mercury  —  Mercuric  Fulminate  —  Sulphates  of  Mer- 
cury—  Mercuric  Sulphide  —  Compounds  of  Mercury  Salts  with 
Ammonia  —  Physiological  Properties  of  Mercury  Compounds  — 
Tests  for  Mercury  —  General  Remarks 393 

CHAPTER  XXIV 
SOLUTIONS,  ELECTROLYSIS,  AND  ELECTRO-CHEMICAL  THEORIES 

Nature  and  Kinds  of  Solutions  —  Absorption  of  Gases  by  Liquids  — 
Solutions  of  Liquids  in  Liquids  —  Solutions  of  Solids  in  Liquids 

—  Degrees  of  Saturation  —  Solid  Solutions — Precipitation  —  Col- 
loidal Solutions  —  Boiling  Points  of  Solutions  —  Use   of  Boiling 
Points  of  Solutions  in  Molecular  Weight  Determinations  —  The 
Freezing  Points  of  Solutions — Discussion  of  Molecular  Weights 


CONTENTS  xv 


Determined  in  Solutions  —  Osmosis  and  Osmotic  Pressure — Elec- 
trolysis —  Electrolytic  Theories  —  Electric  Batteries  —  Electro- 
chemical Series  of  the  Metals 410 

CHAPTER  XXV 
COPPER,  SILVER,  AND  GOLD 

Occurrence,  Metallurgy,  and  Properties  of  Copper  —  Alloys  of  Copper 

—  Oxides  of  Copper  —  Halides  of  Copper — Cyanides  of  Copper 

—  Copper  Salts  of  Oxy-acids  —  Sulphides  of  Copper  —  Analytical 
Tests  for  Copper  —  Occurrence,  Metallurgy,  and  Properties   of 
Silver  —  Oxides  of   Silver  —  Halides   of   Silver — Uses  of   Silver 
Halides  in  Photography  —  Silver  Nitrate  —  Silver  Nitrite  —  Silver 
Sulphate  —  Silver   Carbonate  —  Silver   Phosphate  —  Silver   Sul- 
phide—  Silver   Cyanide  —  Silver  Plating — Silver    Fulminate  — 
Analytical  Tests  for  Silver  —  Occurrence,  Metallurgy,  and  Proper- 
ties of  Gold  —  Gold  Alloys  —  Compounds  of  Gold  —  Analytical 
Tests  for  Gold 436 

CHAPTER  XXVI 
THE  METALS  OF  THE  EARTHS 

Occurrence,  Preparation,  and  Properties  of  Aluminum  —  Uses  of  Alu- 
minum—  Aluminum  Oxide  —  Aluminum  Hydroxide  —  Aluminum 
Chloride  —  Aluminum  Sulphide  —  Aluminum  Sulphate  —  Alums 

—  Aluminum  Silicates  —  Analytical  Tests  for  Aluminum — Gal- 
lium —  Indium  —  Thallium  and  its  Compounds  —  The  Rare-Earth 
Elements 457 

CHAPTER  XXVII 
LEAD  AND  TIN 

Germanium  —  Occurrence,  Metallurgy,  and  Properties  of  Tin  —  Uses 
of  Tin  —  Chlorides  of  Tin  —  Oxides  of  Tin  —  Sulphides  of  Tin  — 
Analytical  Tests  for  Tin  —  Occurrence,  Metallurgy,  and  Proper- 
ties of  Lead  —  Uses  of  Lead  —  Oxides  of  Lead  —  Halides  of  Lead 

—  Lead  Nitrate  —  Lead  Acetate  —  Lead  Sulphate  —  Lead  Arse- 
nate  —  Lead  Carbonate  —  Analytical  Tests  for  Lead      .        .        .    472 

CHAPTER  XXVIII 
CHROMIUM,  MOLYBDENUM,  TUNGSTEN,  AND  URANIUM 

Occurrence,  Preparation,  arid  Properties  of  Chromium  —  Chromic 
Oxides  and  Hydroxides  —  Chromous  Compounds  —  Chromic  Salts 


xvi  CONTENTS 

—  Chromates,  Bichromates,  and  Chromium  Trioxide  —  Chromyl 
Chloride  —  Analytical    Tests    for    Chromium  —  Molybdenum  — 
Tungsten  —  Uranium 485 

CHAPTER  XXIX 
MANGANESE 

Occurrence,  Preparation,  and  Properties  —  Oxides  —  Salts  of  Manga- 
nese—  Manganates  and  Permanganates  —  Uses  of  Permanganates 

—  Analytical  Tests  for  Manganese 495 

CHAPTER  XXX 
IRON,  NICKEL,  AND  COBALT 

Occurrence  of  Iron — Metallurgy  of  Iron  —  Cast  Iron  —  Wrought  Iron 

—  Steel  —  Properties  of  Iron  —  Oxides  and  Hydroxides  of  Iron  — 
Chlorides  of  Iron  —  Sulphides  and  Sulphates  of  Iron  —  Ferrous 
Carbonate  —  Cyanides  of   Iron  —  Blue   Printing  —  Other  Com- 
pounds of  Iron  —  Analytical  Tests  for  Iron  —  Occurrence,  Prepa- 
ration, and  Properties  of  Nickel  —  Nickel  Oxides  and  Hydroxides 

—  Salts  of  Nickel — Nickel  Carbonyl  —  Occurrence,  Preparation, 
and  Properties  of  Cobalt  —  Oxides  and  Hydroxides  of  Cobalt  — 
Other  Cobalt  Compounds  —  Analytical  Tests  for  Cobalt  and  Nickel    502 

CHAPTER  XXXI 
THE  METALS  OF  THE  PLATINUM  FAMILY 

Occurrence  —  Extraction  of  Platinum  from  the  Ores  —  Ruthenium  — 
Rhodium  —  Palladium  —  Osmium  —  Iridium  —  Platinum  —  Ana- 
lytical Tests  for  Platinum 522 

INDEX  .    529 


LIST  OF  ILLUSTRATIONS 


1.  Tube  used  in  demonstrating  that  weight  remains  constant  during 

chemical  changes 10 

2.  Electrolysis  of  water 14 

3.  Preparation  of  hydrogen  by  action  of  sodium  on  water  ...  15 

4.  Preparation  of  hydrogen  by  action  of  steam  on  heated  iron   .         .  15 

5.  Preparation  of  hydrogen  by  action  of  sulphuric  acid  on  zinc  .         .  16 

6.  Transferring  hydrogen  from  one  jar  to  another       ....  17 

7.  Diffusion  of  hydrogen        .........  18 

8.  Formation  of  water  when  hydrogen  burns  in  the  air       .         .         .19 

9.  Singing  flame 20 

10.  A  candle  will  not  burn  in  hydrogen          ......  20 

11.  Oxidation  of  copper  when  heated  in  the  air     .....  21 

12.  Reduction  of  hot  copper  oxide  by  hydrogen 21 

13.  Cylinder  for  compressed  gases 22 

14.  Apparatus  for  determining  hydrogen  equivalents  of  metals    .        .  22 

15.  Burning  of  an  iron  wire  in  oxygen  .......  26 

16.  Burning  of  sulphur  in  oxygen 26 

17.  Oxyhydrogen  blowpipe 31 

18.  Combustion  of  oxygen  in  hydrogen 32 

19.  Lavoisier's  apparatus  to  show  that  mercury  unites  with  oxygen 

when  calcined        ..........  34 

20.  Distillation 37 

21.  Demonstration  of  volumetric  relations  between  oxygen,  hydrogen, 

and  steam 41 

22.  Desiccator 46 

23.  Composition  of  hydrochloric  acid  gas 50 

24.  Electrolysis  of  hydrochloric  acid 51 

25.  Synthesis  of  hydrochloric  acid  gas  by  volume  ....  52 

26.  Burning  of  arsenic  in  chlorine 54 

27.  Action  of  chlorine  on  water  in  sunlight 55 

28.  Ozone  apparatus 88 

29.  Preparation  of  fluorine      .........  103 

30.  Preparation  of  hydrobromic  acid 108 

31.  Sublimation  of  iodine  in  the  laboratory 112 

32.  Sublimation  of  iodine  on  commercial  scale      .....  113 

33.  Titration 129 

34.  Preparation  of  nitrogen  from  the  air 140 

35.  Oxidation  of  nitrogen  by  means  of  the  electric  spark      .         .        .  143 

xvii 


xviil  LIST   OF  ILLUSTRATIONS 


PAGB 


FIG. 

36.  Volumetric  composition  of  ammonia  gas 152 

37.  Decomposition  of  ammonia  by  the  electric  spark  ....  153 

38.  Burning  ammonia  mixed  with  oxygen 154 

39.  Oxidation  of  ammonia  by  use  of  a  platinum  spiral        .        .         .  154 

40.  Preparation  of  nitric  acid 162 

41.  Heating  sodium  in  nitric  oxide       .......  167 

42.  Commercial  distillation  of  sulphur         ......  177 

43.  Crystal  of  rhombic  sulphur 178 

44  to  54.   Crystals  of  the  isometric  system 181 

55  to  60.   Crystals  of  the  tetragonal  system 182 

61  to  67.   Crystals  of  the  hexagonal  system 183 

68  to  71.   Crystals  of  the  orthorhombic  system 184 

72  to  74.   Crystals  of  the  monoclinic  system 184 

75  to  76.    Crystals  of  the  triclinic  system 185 

77.  When  sulphur  burns  in  oxygen  the  volume  remains  unchanged  .  190 

78.  Bleaching  of  flowers  by  means  of  sulphur  dioxide          .         .         .191 

79.  Sulphuric  acid  by  the  contact  process 193 

80.  Diagram  of  a  sulphuric  acid  factory 197 

81  and  82.   Crystal  forms  of  diamond 210 

83.  Acheson  graphite  furnace -      .         .  213 

84.  Typical  arc  furnace  for  experimental  work 213 

85.  Absorption  of  ammonia  gas  by  charcoal 214 

86.  Kipp  apparatus 219 

87.  Siphoning  carbon  dioxide  from  one  jar  to  another         .         .         .221 

88.  Taylor's  carbon  bisulphide  furnace 228 

89.  Yeast  cells        .                 238 

90.  Acetic  acid  organisms 243 

91.  Lactic  acid  organisms 246 

92.  Formulae  of  dextro  and  laevo  lactic  acids 248 

93.  Polariscope 249 

94.  Crystals  of  dextro  and  laevo  tartaric  acid 251 

95.  Grains  of  potato  starch '       .        .  260 

96.  Grains  of  wheat  starch 260 

97.  Grains  of  corn  starch 260 

98.  Potato  starch  grains  in  polarized  light 260 

99.  Manufacture  of  coal  gas 266 

100.    Gases  burn  in  the  flame  of  a  candle 268 

101  to  103.   Demonstration  of  the  reverse  flame 269 

104.  Burning  oxygen  in  coal  gas 270 

105.  Enriching  carbon  monoxide  gas 270 

106.  Principle  of  the  Bunsen  burner 271 

107.  Zones  of  the  flame  of  a  candle        .         .         .                 .        .        .  273 

108  and  109.    A  flame  will  not  pass  through  a  wire  gauze      .        .         .  273 

110.  Davy  safety  lamp 274 

111.  Calorimetric  apparatus 276 

112.  Combustion  bomb  and  calorimeter 277 

113.  Right  and  left  quartz  crystals 292 


LIST   OF   ILLUSTRATIONS  xix 

FIG.  PAGE 

114.  Crystal  of  tridymite 292 

115.  Dialyser 294 

116.  Making  hydrofluosilicic  acid 298 

117.  Retorts  for  making  phosphorus 305 

118.  Electric  furnace  for  making  phosphorus 306 

119.  Making  phosphine  from  phosphorus  and  caustic  alkali          .         .  309 

120.  Phosphine  from  calcium  phosphide 310 

121.  Marsh  test  for  arsenic 319 

122.  Curve  of  atomic  weights  and  atomic  volumes.     (L.  Meyer.)          .  341 

123.  Hopper-shaped  crystal  of  sodium  chloride     .....  356 

124.  Acker  process  of  making  caustic  soda 357 

125.  Salt  cake  furnace,  Le  Blanc  soda  process 358 

126.  Revolving  black  ash  furnace 359 

127.  Solubility  curve  of  sodium  sulphate        ......  362 

128.  Spectroscope 366 

129.  Spectra  of  some  common  elements 367 

130.  Tube  for  examining  spectra  of  gases 368 

131.  Spectra  of  gases 369 

132.  Absorption  of  spectrum  of  blood 370 

133.  Making  metallic  calcium 375 

134.  Common  limekiln 377 

135.  Glass  pots,  open  and  closed  form 384 

136.  Solubility  curve  of  magnesium  chloride          .....  395 

137.  Iron  flask  for  shipping  mercury 402 

138.  Solubility  curves  of  various  salts 413 

139.  Making  colloidal  silver .  417 

140  and  141.   Explanation  of  osmosis 419 

142.  Demonstration  of  osmotic  pressure         ......  421 

143.  Simple  osmometer 421 

144.  Pfeft'er's  osmotic  apparatus 422 

145.  Grotthus's  theory  of  electrolysis 427 

146.  Electrolysis  according  to  the  theory  of  electrolytic  dissociation    .  429 

147.  An  electric  battery 433 

148.  Measuring  the  voltage  of  a  cell 433 

149.  Gravity  battery 434 

150.  Electrolytic  production  of  aluminum 458 

151.  Blastfurnace 503 

152.  Bessemer  converter 507 

153.  Solubility  curve  of  ferric  chloride 511 

154.  Pyrite  crystal 512 

155.  Dobereiner's  lamp 526 


OUTLINES   OF   CHEMISTRY 


CHAPTER  I 

THE  SCOPE  OP  CHEMISTRY  AND  ITS  RELATIONS  TO 
OTHER  SCIENCES 

CHEMICAL  CHANGE,  ELEMENTS,  AND  COMPOUNDS 

OUR  own  bodies,  and  the  various  objects  that  surround  us, 
constitute  the  subject  of  study  of  the  natural  sciences.  The 
investigation  of  the  things  that  make  up  the  universe  as  we 
know  it,  is  conducted  by  means  of  our  senses,  either  aided  or 
unaided.  For  the  sake  of  classifying  our  knowledge,  we  are 
wont  to  distinguish  between  the  biological  sciences,  which  deal 
with  living  things,  and  the  so-called  physical  sciences  of  astron- 
omy, geology,  physics,  and  chemistry.  Astronomy,  which 
deals  with  the  heavenly  bodies,  is  nevertheless  closely  related 
to  the  sciences  of  physics  and  chemistry,  though  obviously  not 
to  biology.  But  the  study  of  living  things  and  the  life  history 
of  the  earth  and  the  processes  that  are  continually  going  on  on 
its  surface  is  inseparably  linked  with  the  subjects  of  physics 
and  chemistry.  The  latter  sciences  may  indeed  be  regarded  as 
basal  in  character.  The  study  of  matter  —  that  is,  anything 
which  occupies  space  —  comes  within  the  scope  of  these  two 
sciences.  Viewed  in  this  light,  biology,  astronomy,  and  geol- 
ogy merely  present  special  complex  phases  and  combinations  of 
physics  and  chemistry. 

Physical  and  Chemical  Changes. — The  changes  which  any 
object  may  undergo  are  either  superficial  or  deep-seated  in 
character.  Thus,  if  a  stick  of  sulphur  be  thrown  or  whirled 
through  the  air,  the  character  of  the  sulphur  is  not  altered, 
though  the  sulphur  has  undergone  change  of  position  through 
expenditure  of  mechanical  energy  upon  it.  Energy  is  any- 
thing which  does  work  or  is  capable  of  doing  work.  Energy 
itself  is  measured  by  the  amount  of  work  it  has  done  or  is 


2  OUTLINES   OF  CHEMISTRY 

capable  of  doing.  Indeed,  as  energy  is  always  measured  in 
terms  of  work,  the  two  are  often  regarded  as  synonymous. 
Work  is  equal  to  the  force  multiplied  by  the  distance  through 
which  the  force  acts,  a  force  being  defined  as  that  which  causes 
or  modifies  motion,  the  latter  being  a  change  of  place.  The 
motion  might  have  been  imparted  to  the  sulphur  by  means  of  the 
muscles  or  by  a  contrivance  in  which  the  energy  was  furnished 
by  gravity,  heat,  light,  electricity,  magnetism,  etc.  These 
agencies  are  consequently  capable  of  doing  work ;  that  is,  they 
represent  forms  of  energy.  As  long  as  the  sulphur  remains 
sulphur,  no  matter  through  what  motions  or  other  alterations, 
like  contraction,  expansion,  electrification,  change  of  temper- 
ature, pulverization,  liquefaction,  or  vaporization,  it  may  go, 
the  change  in  question  is  called  a  physical  change,  and  the 
study  of  such  changes  in  all  their  various  phases  belongs  to  the 
subject  of  physics.  But  if,  for  example,  we  burn  the  sulphur 
in  the  air  we  obtain  a  gas  of  a  pungent  odor  which  may  be  con- 
densed with  the  aid  of  pressure  and  lowering  of  the  tempera- 
ture to  a  colorless,  mobile  liquid.  This  is  quite  unlike  sulphur 
in  all  its  various  properties,  and  we  consequently  say  that  a 
new  substance  has  been  formed.  The  process  of  forming  a 
new  substance  is  called  a  chemical  change. 

Any  process  in  which  given  substances  disappear  and  new  ones 
are  formed  is  chemical,  and  the  study  of  such  deep-seated  processes 
in  all  their  various  phases  is  the  subject  with  which  the  science  of 
chemistry  is  concerned.  It  would  thus  seem  fairly  easy  to  dis- 
tinguish between  chemical  and  physical  processes.  Indeed,  in 
general,  such  a  distinction  can  readily  be  made  on  the  basis  of 
what  has  just  been  said.  But  whether  new  substances  have 
been  formed  must  be  decided  from  the  properties  of  the  material; 
and  there  must  consequently  be  some  definite  way  of  telling 
whether  an  alteration  of  substance  has  occurred  or  not.  It  is 
evident  at  once  that  the  term  substance  must  be  clearly  defined. 
For  our  present  purpose,  it  will  suffice  to  say  that  a  substance 
is  matter  which  is  perfectly  homogeneous  throughout,  considered 
without  respect  to  shape  or  amount.  Thus  sulphur,  iron,  and 
water  are  substances.  Many  things  which  are  apparently 
homogeneous  in  character  are  not  so  in  reality.  Thus,  the 
atmosphere  on  closer  study  is  found  to  be  a  mixture  of  nitrogen, 
oxygen,  carbon  dioxide,  and  other  gases ;  sea  water  is  found  to 


THE   SCOPE   OF   CHEMISTRY  3 

consist  of  water  together  with  various  saline  substances ;  brass 
is  made  up  of  copper  and  zinc  in  proportions  that  may  vary  to 
a  considerable  extent  in  different  samples. 

If  we  pulverize  a  piece  of  roll  sulphur  and  grind  it  together 
with  iron  filings  in  a  mortar  as  intimately  as  possible,  a  fine 
grayish  powder  results  which  has  the  outward  appearance  of 
homogeneity.  On  closer  inspection,  however,  with  the  aid  of  a 
microscope  perchance,  this  powder  appears  heterogeneous;  in 
other  words,  it  is  merely  a  physical  mixture.  Indeed,  it  is  very 
easy  to  separate  the  iron  from  the  sulphur,  for  by  passing  a 
magnet  through  the  mixture  the  iron  will  adhere  to  the  magnet, 
and  the  sulphur  will  be  left  behind.  We  could  also  separate 
the  iron  from  the  sulphur  in  the  mixture  by  treating  the  latter 
with  carbon  disulphide,  which  liquid  dissolves  the  sulphur  and 
leaves  the  iron  unaltered.  The  mixture  of  iron  filings  and  sul- 
phur represents  a  typical  physical  mixture.  It  is  obviously 
heterogeneous  in  character,  the  proportion  of  iron  and  sulphur 
in  the  mixture  may  be  varied  at  will,  and  the  iron  and  sulphur 
may  readily  be  separated  from  each  other  by  simple  means. 

If  now  we  heat  some  of  the  mixture  of  pulverized  sulphur 
and  iron  filings  in  a  test  tube,  we  observe  that  at  a  certain 
temperature  the  contents  of  the  tube  begin  to  glow.  As  we 
take  it  out  of  the  flame  the  glowing  nevertheless  increases,  and 
the  contents  of  the  tube  become  hotter.  After  a  time  the  glow- 
ing becomes  weaker,  and  gradually  ceases  as  the  material  cools. 
It  is  evident  that  by  raising  tlie  temperature  of  the  mixture  of 
iron  filings  and  sulphur  to  a  certain  point,  a  change  was  inaugu- 
rated, which  on  taking  away  the  source  of  heat  nevertheless 
continued,  giving  off  additional  heat  and  light.  On  examining 
the  contents  of  the  tube  after  it  has  cooled  to  room  temperature, 
we  find  a  black  mass,  quite  unlike  either  the  sulphur  or  iron  in 
appearance.  We  can  no  longer  detect  heterogeneity  in  it  even 
with  the  aid  of  the  microscope.  The  magnet  is  unable  to 
extract  iron  from  this  material,  and  carbon  disulphide  will  not 
alter  it  in  any  way.  A  few  drops  of  hydrochloric  acid  poured 
upon  it  evolve  a  malodorous  gas  called  hydrogen  sulphide,  which 
is  not  formed  when  a  simple  mixture  of  iron  filings  and  sulphur 
is  moistened  with  that  acid.  We  clearly  have  formed  a  new 
substance  by  heating  the  sulphur  and  iron  together.  It  is  called 
ferrous  sulphide,  and  results  from  simple  union  of  sulphur  and 


4  OUTLINES  OF   CHEMISTRY 

iron  at  elevated  temperatures.  It  has  been  found  that  ferrous 
sulphide  contains  63.52  per  cent  of  iron  and  36.48  per  cent  of 
sulphur,  and  that  it  always  has  exactly  this  composition  no 
matter  by  what  methods  it  has  been  formed.  This  is  in  fact  a 
characteristic  of  all  chemical  compounds.  We  may  express 
this  fact  by  saying  that  every  definite  chemical  compound  always 
contains  the  same  ingredients  in  the  same  proportion  by  weight. 
This  is  the  law  of  definite  proportions.  A  law,  as  the  word  is 
used  in  science,  is  a  general  statement  summarizing  what  has  actu- 
ally been  found  to  be  true  in  a  large  number  of  individual  cases 
that  have  been  carefully  investigated. 

Other  typical  examples  of  chemical  change  are  the  rusting  of 
iron,  the  combustion  of  coal  or  wood,  the  decomposition  of 
water  by  electrolysis,  the  formation  of  quicklime  from  lime- 
stone by  the  agency  of  heat,  the  change  of  carbon  dioxide  and 
water  into  starch  by  sunlight  in  the  green  leaf  of  the  plant, 
and  the  darkening  of  a  photographic  plate  when  exposed  to 
light.  In  all  these  cases  new  substances  are  formed,  and  the 
actions  are  accompanied  by  changes  in  temperature,  volume, 
outward  appearance,  and  other  specific  properties  which  char- 
acterized the  original  substances  before  the  change  occurred. 
It  is  the  province  of  chemistry  to  study  such  changes  in  all 
their  various  aspects.  This  involves  a  close  study  of  the  com- 
position and  specific  properties  of  the  substances  before  and 
after  the  chemical  change,  which  is  commonly  termed  the  chem- 
ical reaction,  has  taken  place.  But,  in  addition,  a  study  of  the 
conditions  that  must  obtain  in  order  that  the  reaction  may 
begin  and  proceed,  and  an  investigation  of  the  various  energy 
changes  that  accompany  the  reaction,  also  fall  within  the  field 
of  chemistry.  Thus  we  have  various  branches  of  chemistry. 
So  analytical  chemistry  seeks  to  determine  the  qualitative  arid 
quantitative  composition  of  substances  by  tearing  them  apart  or 
analyzing  them ;  synthetic  chemistry  seeks  to  build  up  more 
complex  substances  from  simpler  ones  ;  thermochemistry  concerns 
itself  with  the  thermal  changes  accompanying  chemical  reac- 
tions ;  electrochemistry  is  concerned  with  electricity  as  an  agent 
in  producing  chemical  changes,  or  as  an  accompaniment  of 
chemical  phenomena ;  photochemistry  treats  of  the  relations  of 
light  to  chemical  changes.  In  the  crust  of  the  earth,  in  the 
atmosphere,  in  natural  waters,  in  the  bodies  of  plants  and  ani- 


THE   SCOPE   OF   CHEMISTRY  5 

mals,  chemical  changes  are  continually  going  on.  Upon  these 
all  life  on  the  globe  depends.  Every  breath  we  breathe,  every 
move  we  make,  every  thought  we  think,  is  accompanied  by 
chemical  changes  and  their  concomitant  physical  phenomena  as 
above  briefly  mentioned.  The  importance  of  the  study  of  chem- 
istry, therefore,  is  clearly  apparent,  and  it  is  also  evident  why 
there  must  needs  be  many  special  and  applied  lines  of  this  sub- 
ject, which  seek  to  investigate  certain  special  fields.  Thus  we 
have  agricultural  chemistry,  pharmaceutical  chemistry,  physio- 
logical chemistry,  food  chemistry,  industrial  chemistry,  etc.,  the 
province  of  each  of  which  is  indicated  sufficiently  by  the  name 
itself. 

From  what  has  been  stated,  it  would  seem  a  fairly  simple  mat- 
ter to  distinguish  a  chemical  change  from  a  purely  physical  one, 
but  this  is  by  no  means  always  easy.  Suppose  a  block  of  ice  and 
one  of  common  salt  be  placed  in  contact  with  each  other ;  we  note 
that  the  salt  and  ice  gradually  disappear,  forming  a  brine.  Evi- 
dently the  brine  has  quite  different  properties  from  those  of 
either  the  salt  or  the  ice.  Moreover,  there  was  a  marked  change 
of  temperature,  in  this  case  a  cooling  effect,  as  the  salt  and  ice 
acted  on  each  other.  Furthermore,  a  contraction  ensued,  for  the 
volume  of  the  brine  is  less  than  the  sum  of  the  volumes  of  the 
blocks  of  ice  and  salt.  Again,  as  a  block  of  ice  and  one  of  par- 
aifine,  or  one  of  salt  and  one  of  paraifine,  for  example,  do  not 
act  on  each  other  at  all  when  brought  into  contact,  it  is  clear 
that  the  action  between  ice  and  salt  takes  place  because  of  the 
specific  nature  of  the  substances.  Furthermore,  it  has  been 
found  that  below  —  22°  C.  ice  and  common  salt  no  longer  act  on 
each  other,  just  as  iron  and  sulphur  do  not  act  on  each  other  at 
ordinary  temperatures.  Raise  the  temperature  sufficiently  in 
each  case,  and  at  a  certain  definite  point  action  begins.  Thus, 
the  interaction  of  ice  and  common  salt  apparently  bears  all  the 
earmarks  of  a  genuine  chemical  change.  This  is  indeed  true 
except  in  one  particular  which  has  not  yet  been  mentioned, 
namely,  it  is  possible  to  vary  the  composition  of  the  brine  grad- 
ually, by  adding  common  salt  to  it  till  a  point  of  saturation  is 
reached.  Even  then  the  brine  will  still  take  up  somewhat  more 
salt  gradually  if  the  temperature  of  the  whole  is  slowly  raised. 
The  brine  is  termed  a  solution  of  common  salt  in  water.  It 
results  from  the  action  of  salt  and  water  on  each  other.  The 


6  OUTLINES  OF  CHEMISTRY 

water  used  may  be  liquid,  or  in  form  of  ice  above  —22°  C.  A 
distinction  is  commonly  made  between  solutions  and  chemical 
compounds.  In  a  solution,  the  relative  amounts  of  the  ingredients 
that  it  contains  may  be  varied  gradually  within  certain  limits,  as 
we  have  seen  in  the  case  of  the  brine.  In  a  chemical  compound, 
the  constituents  cannot  thus  be  varied  in  amount.  Not  many 
years  ago,  chemists  spoke  of  solutions  as  chemical  combinations 
according  to  variable  proportions,  and  this  term  is  indeed  indic- 
ative of  the  real  relation  that  they  bear  to  definite  chemical 
compounds  which  follow  the  law  of  definite  proportions. 

Brine,  then,  is  not  a  mere  physical  mixture,  and  it  is  conse- 
quently not  to  be  classed  with  such  mixtures  as  that  of  sulphur 
and  iron  filings  rubbed  together  in  a  mortar,  which  represents 
a  typical  physical  mixture.  In  chemistry  we  frequently  have 
to  deal  with  (1)  physical  mixtures,  (2)  solutions  (i.e.  com- 
pounds according  to  variable  proportions),  and  (3)  definite 
chemical  compounds. 

As  further  typical  examples  of  solutions  may  be  mentioned, 
solution  of  sugar  in  water,  of  camphor  in  petroleum  oil,  of 
ether  in  alcohol,  of  carbon  disulphide  in  olive  oil.  The  subject 
of  solutions  clearly  forms  an  important  part  of  chemistry,  and 
it  will  consequently  be  considered  more  fully  later. 

Chemical  Elements.  —  A  careful  study  of  all  substances 
known  has  revealed  the  fact  that  there  are  about  eighty  which 
it  has  been  impossible  to  decompose  into  simpler  substances 
thus  far.  These  substances  are  regarded  as  elementary  in  char- 
acter. They  are  termed  the  chemical  elements.  Whether  a 
substance  is  an  element  or  not  is  thus  determined  by  experi- 
ment. As  new  methods  of  experimental  attack  are  discovered, 
substances  that  are  now  regarded  as  elements  may  prove  to  be 
complex  and  consequently  capable  of  synthesis.  Thus  at  one 
time  lime  and  caustic  potash  were  regarded  as  elements,  whereas 
now  we  know  that  lime  contains  calcium  and  oxygen,  and 
caustic  potash  consists  of  potassium,  hydrogen,  and  oxygen. 
Sir  William  Ramsay  found  that  the  emanations  from  radium 
show  the  spectra  of  helium,  argon,  and  neon,  and  this  is  by 
many  regarded  as  a  case  of  synthesis  of  the  latter  gases  from 
the  products  of  the  decay  of  radium.  Again,  Ramsay  claims 
to  have  obtained  spectroscopic  traces  of  lithium  by  the  action 
of  radium  emanation  upon  copper  sulphate  solutions,  though 


THE   SCOPE   OF   CHEMISTRY  7 

Mme.  Curie's  investigations  do  not  substantiate  his  results. 
Thus  it  is  evident  that  some  of  the  substances  we  now  term 
elements  may  prove  to  be  composite.  It  is  also  obviously  im- 
possible to  state  just  how  many  elements  there  are,  for  it  is 
uncertain  whether  some  substances  are  elementary  or  complex 
in  character.  The  following  is  an  alphabetical  list  of  the 
chemical  elements  as  commonly  recognized  at  present. 

CHEMICAL  ELEMENTS 

Aluminum  Europium  Mercury  Silicon 

Antimony  Fluorine  Molybdenum  Silver 

Argon  Gadolinium  Neodymium  Sodium 

Arsenic  Gallium  Neon  Strontium 

Barium  Germanium  Nickel  Sulphur 

Bismuth  Glucinum  Nitrogen  Tantalum 

Boron  Gold  Osmium  Tellurium 

Bromine  Helium  Oxygen  Terbium 

Cadmium  Hydrogen  Palladium  Thallium 

Caesium  Indium  Phosphorus  Thorium 

Calcium  Iodine  Platinum  Thulium 

Carbon  Iridium  Potassium  Tin 

Cerium  Iron  Praseodymium  Titanium 

Chlorine  Krypton  Radium  Tungsten 

Chromium  Lanthanum  Rhodium  Vanadium 

Cobalt  Lead  Rubidium  Xenon 

Columbium  Lithium  Ruthenium  Ytterbium 

Copper  Lutecium  Samarium  Yttrium 

Dysprosium  Magnesium  Scandium  Zinc 

Erbium  Manganese  Selenium  Zirconium 

It  will  be  observed  that  the  list  contains  a  goodly  number 
of  common,  well-known  substances.  Notably,  it  appears  that 
all  the  metals  are  elements.  Again,  there  are  substances  in 
the  list  which  are  not  metals,  like  sulphur,  chlorine,  bromine, 
iodine,  oxygen,  hydrogen,  phosphorus,  etc.  The  elements  may  be 
divided  into  two  groups  ;  namely,  the  metals  and  non-metals.  It 
is  difficult  to  draw  a  sharp  line  between  these  groups,  however, 
for  elements  like  arsenic,  antimony,  and  tellurium  clearly  rep- 
resent transitions  between  the  metals  and  non-metals.  Such 
transition  elements  are  sometimes  called  metalloids. 

Some  of  the  elements  are  gases,  others  are  liquids,  and  still 
others  are  solids,  under  ordinary  conditions  of  temperature 
and  pressure.  Whether  an  element  is  a  solid,  liquid,  or  gas 


8 


OUTLINES   OF   CHEMISTRY 


is  determined  entirely  by  the  conditions  of  temperature  and 
pressure  to  which  it  is  subjected. 

Less  than  half  of  the  elements  mentioned  in  the  table  enter 
into  the  composition  of  ordinary  objects.  The  solid  crust  of 
the  earth,  also  called  the  lithosphere,  makes  up  about  93  per 
cent  of  all  known  terrestrial  matter,  while  the  ocean  represents 
about  7  per  cent,  and  the  atmosphere  only  0.03  per  cent,  of  the 
total.  The  following  table,  by  F.  W.  Clarke,  gives  an  estimate 
of  the  relative  amounts  of  the  elements  contained  in  the  litho- 
sphere and  the  ocean.  The  third  column  of  the  table  gives  a 
total  average  including  the  atmosphere. 

AVERAGE   COMPOSITION   OF   LITHOSPHERE,   OCEAN,  AND 
ATMOSPHERE 


LITHOSPHERE 
(93  PEE  CENT) 

OCEAN 
(7  PER  CENT) 

AVERAGE 
INCLUDING  THE 
ATMOSPHERE 

Oxygen  
Silicon  ...... 

47.07 

28.06 
7.90 

85.79 

49.78 
26.08 
7.34 

Iron  • 

443 

411 

3.44 

0.05 

3.19 

Magnesium  
Sodium  
Potassium  

2.40 
2.43 
2.45 
0.22 

0.14 
1.14 
0.04 
10.67 

2.24 
2.33 

2.28 
0.95 

0.40 

0.37 

Carbon  ...... 
Chlorine  

0.20 
0.07 

0.002 
2.07 
0.008 

0.19 
0.21 

0.11 

0.11 

Sulphur  

0.11 
0.09 

0.09 

0.11 
0.09 

0.07 

0.07 

0.03 

0.03 

Nitrogen  ...... 
Fluorine  ...... 
All  other  elements  .... 

0.02 
0.50 



0.02 
0.02 
0.48 

100.00 

100.00 

100.00 

The  following  table  gives  the  approximate  amounts  of  the 
elements  found  in  the  human  body :  — 


THE   SCOPE   OF   CHEMISTRY  9 

AVERAGE    ELEMENTARY  COMPOSITION  OF  THE   HUMAN 

BODY 

Oxygen 66.0  per  cent 

Carbon 17.6  per  cent 

Hydrogen 10.1  per  cent 

Nitrogen       . 2.5  percent 

Calcium 1.5  per  cent 

Phosphorus 1.0  percent 

Potassium 0.4  per  cent 

Sodium 0.3  per  cent 

Chlorine 0.3  per  cent 

Sulphur 0.25  per  cent 

Magnesium  .        .        .        .        .  .        .0.04  per  cent 

Iron 0.004  per  cent 

Silicon,  Fluorine,  Iodine,  etc.,  in  traces. 

Compounds.  —  Most  substances  are  non-elementary  in  charac- 
ter ;  that  is,  they  are  combinations  of  two  or  more  elements. 
Such  substances  are  consequently  termed  compounds.  They 
may  be  formed  by  direct  union  of  the  elements  with  one 
another  under  proper  conditions ;  as,  for  instance,  sulphur  may 
unite  with  iron  to  form  ferrous  sulphide.  Again,  limestone, 
which  is  carbonate  of  calcium,  decomposes  at  a  high  tempera- 
ture, forming  two  simple  substances,  lime  and  carbon  dioxide, 
the  latter  being  a  gas.  Further,  by  action  of  two  compounds 
on  each  other,  two  other  compounds  may  result.  As  an 
example,  when  common  salt  and  nitrate  of  silver  are  brought 
together  in  aqueous  solution,  silver  chloride,  a  substance  in- 
soluble in  water,  and  sodium  nitrate,  a  soluble  substance,  are 
formed.  This  latter  change  is  termed  double  decomposition 
or  metathesis. 

In  the  three  cases  cited  we  have,  indeed,  the  three  types  of 
chemical  changes ;  namely,  (1)  the  direct  union  of  two  or  more 
substances  to  form  a  single  compound,  (2)  the  breaking  up  of  a 
compound  into  simpler  ones,  and  (3)  the  interaction  of  substances 
with  one  another  to  form  new  substances. 

Like  elements,  compounds  may  also  assume  the  solid,  liquid, 
or  gaseous  state,  according  to  the  conditions  of  temperature  and 
pressure  that  obtain.  However,  by  no  means  all  compounds  are 
capable  of  assuming  these  three  states,  for  many  readily  decom- 
pose when  an  attempt  is  made  to  liquefy  them  or  volatilize  them 
by  means  of  heat. 


10 


OUTLINES  OF  CHEMISTRY 


Compounds  which  contain  different  elements  are,  of  course, 
different  in  character.  The  same  is  true  of  compounds  that 
contain  the  same  elements,  though  in  different  proportions  by 
weight.  For  a  long  time  it  was  thought  that  one  compound 
could  differ  from  another  only  because  it  contained  either  differ- 
ent elements,  or  the  same  elements  in  different  proportions. 
However,  we  now  have  knowledge  of  a  large  number  of  com- 
pounds that  are  quite  different  substances,  and  yet  they  con- 
tain the  same  elements  in  exactly  the  same  proportion  by 
weight.  Such  compounds  are  called  isomers,  and  the  difference 
between  them  is  explained  by  the  different  manner  in  which  the 
elements  are  combined  in  these  substances,  in  other  words,  by  the 
difference  in  inner  structure  or  constitution  of  the  compounds. 
Conservation  of  Mass.  —  Investigations  have  shown  that  when 
chemical  changes  take  place,  the  weight  of  all  the  substances 
before  the  reaction  is  equal  to  the  weight  of  all  the  substances 
after  the  reaction  has  taken  place.  In  other  words,  in  any 
chemical  change  the  total  weight  remains  the  same.  As  at  any 
place  on  the  surface  of  the  earth  weight  and  mass  are  propor- 
tional to  each  other,  we  may  say  that  during  chemical  changes, 
the  total  mass  of  the  reacting  substances  remains  constant.  This 
is  simply  the  law  of  conservation  of  mass,  which  applies  to  chemi- 
cal as  well  as  to  physical  changes.  It  is  sometimes  called  the 
law  of  conservation  of  matter.  It  is  the  outcome  of  experi- 
mental investigations,  the  most  careful  of  which  were  conducted 
by  having  chemical  changes  go  on  in  sealed  glass  vessels,  which, 

together  with  their  contents,  were 
weighed  before  and  after  the  sub- 
stances they  contained  had  reacted 
chemically  on  one  another. 

Figure  1  shows  a  common  type  of 
such  sealed  glass  tubes.  The  sub- 
stances are  introduced  into  each  limb, 
and  the  tube  is  then  sealed  by  drawing 
off  the  end  as  shown.  After  the  whole 
has  been  very  accurately  weighed,  the 
contents  are  allowed  to  act  on  each 
other  by  inclining  or  shaking  the  tube.  After  the  action 
has  ceased  and  the  whole  has  cooled  to  room  temperature,  the 
tube  is  carefully  weighed  again.  H.  Landolt  has  performed 


FIG.  l. 


THE   SCOPE   OF   CHEMISTRY  11 

many  careful  experiments  of  this  nature  in  recent  years.  His 
results  show  that  if  there  is  any  change  of  weight,  it  lies  very 
near  the  limit  of  experimental  error.  That  is,  it  is  so  slight  as 
to  be  quite  negligible  for  all  ordinary  purposes. 

Conservation  of  Energy.  —  Like  mass,  energy  also  cannot  be 
created  or  destroyed.  It  can  simply  be  transformed.  Thus,  for 
example,  electricity  may  be  converted  into  heat,  mechanical 
energy,  or  chemical  energy ;  and  again,  each  of  these  latter  may 
be  converted  back  into  electricity. 

When  coal  burns,  to  be  sure,  new  substances  are  being  formed, 
but  in  addition  chemical  energy  is  being  converted  into  heat 
and  light.  When  water  is  decomposed  by  means  of  the  elec- 
tric current,  electrical  energy  is  being  converted  into  chemical 
energy.  When  lime  is  produced  at  the  high  temperature  of  the 
limekiln,  heat  is  transformed  into  chemical  energy.  When 
starch  is  formed  in  the  sunlight  in  the  green  leaf  of  the  plant, 
light  is  converted  into  chemical  energy. 

The  Cause  of  Chemical  Change.  —  As  to  the  cause  why  certain 
substances  act  on  each  other  to  form  new  substances  under 
given  conditions  and  other  substances  do  not,  we  are  quite 
ignorant.  Thus,  we  cannot  tell  why  a  piece  of  sulphur  will 
burn  when  heated  in  the  air  and  a  piece  of  platinum  or  gold 
will  not.  We  know  that,  in  the  act  of  burning,  the  sulphur 
unites  with  the  oxygen  of  the  air,  and  therefore  we  explain  this 
by  saying  that  sulphur  and  oxygen  have  a  specific  attraction 
for  each  other.  This  specific  attraction,  which  is  regarded  as 
the  cause  of  chemical  union,  is  commonly  called  chemical  affinity. 
Thus,  the  fact  that  platinum  and  gold  do  not  burn  when  heated 
in  the  air  is  explained  by  saying  that  these  elements  have  too 
slight  a  chemical  affinity  for  oxygen. 

The  word  affinity  means  relationship.  It  was'  adopted  at  a 
time  when  it  was  thought  that  substances  that  are  similar  are 
more  prone  to  unite  chemically  with  one  another  than  those 
which  are  dissimilar.  While  it  is  true,  as  we  shall  see,  that 
substances  of  similar  characteristics  do  frequently  unite  chem- 
ically, nevertheless,  as  a  rule,  substances  that  are  unlike  in 
character  react  more  energetically  with  one  another.  So,  for 
instance,  while  metals  do  form  chemical  compounds  with  metals, 
yet  they  react  much  more  energetically  with  non-metals  and 
thus  form  stabler  compounds. 


12  OUTLINES   OF   CHEMISTRY 

Factors  affecting  Chemical  Change.  —  In  order  that  a  chemi- 
cal change  may  take  place,  it  is  first  of  all  necessary  that  the 
substances  that  are  brought  together  be  of  the  right  kind  ;  that 
is,  they  must  be  of  such  a  specific  nature  that  they  will  react. 
According  to  the  preceding  paragraph,  we  should  say,  the  sub- 
stances must  have  chemical  affinity  for  one  another.  When  we 
study  any  substance  as  to  its  power  to  react  with  other  bodies 
to  form  new  substances,  we  are  investigating  the  chemical 
properties  of  the  substances.  Intimate  contact  of  the  substances 
that  are  to  react  is  always  necessary.  From  this  fact  we  con- 
clude that  chemical  affinity  acts  at  insensible  distances.  Again, 
temperature  is  a  great  factor  in  promoting  chemical  change ;  in- 
deed, in  most  of  the  changes  studied  in  the  laboratory,  tempera- 
ture is  second  only  to  chemical  affinity  itself  in  determining 
whether  chemical  action  will  take  place  or  not.  Electricity, 
light,  pressure,  concussion,  various  forms  of  vibration,  contact  with 
other  substances  which  often  need  to  be  present  only  in  relatively 
minute  quantity,  are  also  frequently  important  factors  in  determin- 
ing whether  a  chemical  change  will  proceed  or  not.  Furthermore, 
the  relative  amounts  of  the  reacting  substances  brought  into  contact 
also  affect  the  rate  of  a  chemical  change  and  the  extent  or  degree 
of  completion  to  which,  it  will  proceed. 


CHAPTER  II 

HYDROGEN 

History.  —  It  was  known  to  Paracelsus  (1493-1541)  that  an 
inflammable  gas  is  produced  when  dilute  acids  act  on  certain 
metals ;  but  the  English  physicist  Cavendish  (1731-1810)  was 
the  first  to  isolate  hydrogen  and  recognize  it  as  a  special  gas. 
In  1766  he  prepared  hydrogen  by  the  action  of  either  hydro- 
chloric or  sulphuric  acid  on  zinc,  iron,  or  tin,  and  described  the 
characteristic  properties  of  the  gas.  Hydrogen  is  an  essential 
constituent  of  water,  and  derives  its  name  from  the  Greek  words 
meaning  water  and  to  generate. 

Occurrence. —  Hydrogen  is  perhaps  the  most  widely  distributed 
element  in  the  universe.  It  occurs  in  very  large  quantities  in 
the  sun,  where  it  is  heated  to  incandescence  owing  to  the  high 
temperature  that  obtains.  It  is  found  in  all  fixed  stars  and 
nebulae  that  have  been  examined  by  means  of  the  spectroscope. 
On  the  earth  it  occurs  only  in  small  amounts  in  the  free  state. 
The  atmosphere  contains  only  about  0.005  per  cent  of  uncom- 
bined  hydrogen  by  volume.  In  the  gases  emitted  from  vol- 
canoes, oil  wells,  and  some  natural  salt  deposits,  notably  those 
at  Stassfurt,  hydrogen  is  found  in  the  free  state.  It  occurs 
further  in  the  gases  resulting  from  certain  forms  of  fermentation, 
in  the  gases  emitted  from  living  plants,  and  in  the  intestinal 
gases  of  human  beings  and  animals.  In  meteoric  iron,  and  in 
various  minerals,  hydrogen  has  also  at  times  been  found  as  an 
occlusion. 

While  hydrogen  exists  only  in  small  quantities  in  the  free 
state  on  the  earth,  in  combination  with  other  elements  it  is  found 
in  very  large  quantities.  Thus,  11.19  per  cent  of  the  weight  of 
water  consists  of  hydrogen.  It  forms  an  essential  part  of  all 
plants  and  animals,  in  which  it  occurs  chiefly  in  combination 
with  the  elements  oxygen,  carbon,  and  nitrogen.  In  petroleum, 
natural  gas,  and  marsh  gas  it  occurs  combined  with  carbon.  It 
is  an  essential  constituent  of  all  acids. 

13 


14 


OUTLINES  OF  CHEMISTRY 


Preparation. — When  an  electric  current  is  passed  through 
water  acidulated  with  sulphuric  acid  (Fig.  2),  both  hydrogen 

and  oxygen  are  produced, 
two  volumes  of  the  former 
and  one  volume  of  the 
latter  appearing  at  the 
opposite  plates  used  as 
electrodes.  This  method 
is  an  excellent  one  for 
preparing  very  pure  hy- 
drogen. The  process 
itself  is,  however,  some- 
what complex  in  nature 
and  will  receive  special 
attention  later  (see  Elec- 
trolysis). 

When .  metallic  sodium 
acts  on  water,  hydrogen 
and  caustic  soda  are 
formed.  The  sodium 
may  be  introduced  into  a 
test  tube  which  has  been 
filled  with  water  and  in- 
verted in  a  basin  as  shown 
in  Fig.  3.  The  metal, 
being  lighter  than  water, 
rises  in  the  tube,  and  as 
the  hydrogen  is  generated 
it  forces  the  water  out  of 

FlG  2  the     tube.       The     metal 

melts  owing  to  the  heat 

generated  during  the  reaction  and  floats  in  form  of  a  globule 
on  top  of  the  water.  We  may  express  what  takes  place  by 
writing :  — 

Water  +  Sodium  =  Hydrogen  +  Caustic  Soda. 

The  latter  substance  is  dissolved  in  the  water  after  the  change 
has  taken  place.  It  may  be  obtained  as  a  white  solid  by 
boiling  the  solution  till  all  the  water  has  evaporated.  The 
caustic  soda  solution  turns  red  litmus  blue,  has  an  "  alkaline  " 


HYDROGEN 


15 


taste,  and  feels  slippery  to  the  touch.  Caustic  soda  consists 
of  three  elements,  sodium,  oxygen,  and  hydrogen.  It  is  also 
called  sodium  hydroxide.  It  is 
a  strong  alkali,  that  is,  a  sub- 
stance which  is  able  to  neutralize 
acids  and  thus  form  salts. 

Potassium  acts  on  water  like 
sodium,  only  much  more  vigor- 
ously. The  metal  in  this  case 
catches  fire  and  burns  with  a 
brilliant  flame.  Frequently  the 
action  is  so  violent  as  to  result 
in  explosions.  Lithium,  rubid- 
ium, caesium,  barium,  strontium, 
and  calcium  also  act  on  water 
at  room  temperature,  forming 
hydrogen  and  the  hydroxide  of 

the  metal  employed.  It  is  therefore  evident  that  all  of  these 
metals  cannot  be  kept  in  contact  with  the  air,  which  always 
contains  some  moisture.  They  are  kept  under  hydrocarbon 
oils,  like  kerosene,  with  which  they  do  not  react. 

Magnesium  decomposes  water  at  room  temperatures,  but 
very  slowly  indeed.  If,  however,  a  magnesium  salt  is  dis- 
solved in  the  water,  the  action  goes  on  much  more  rapidly. 
Magnesium  salts  aid  the  action  by  dissolving  the  magnesium 
hydroxide  formed,  which  would  inclose  the  metal  in  a  pro- 


FIG.  3. 


FIG.  4, 


16 


OUTLINES   OF  CHEMISTRY 


tecting  film.  On  boiling  water  magnesium  acts  quite  readily, 
forming  hydrogen  and  the  hydroxide  of  the  metal.  Zinc  or 
iron  when  heated  to  redness  in  a  tube  will  decompose  steam, 
yielding  hydrogen  and  an  oxide  of  the  metal  employed  (Fig.  4). 
Furthermore,  by  similarly  passing  steam  over  red-hot  carbon, 
hydrogen  and  carbon  monoxide  are  formed.  This  latter  process 
is  used  in  making  water  gas  (which  see). 

By  boiling  zinc  in  aqueous  caustic  potash  solution,  hydrogen 
and  potassium  zincate  result.  The  latter  is  a  salt  which 
remains  in  solution;  thus:  — 

Caustic  Potash  +  Zinc  =  Potassium  Zincate  4-  Hydrogen. 

Caustic  soda  acts  like  caustic  potash.  Aluminum  acts  in  a 
manner  similar  to  zinc.  In  this  case  an  aluminate  instead  of 
a  zincate  is  formed  and  remains  in  solution.  By  heating  zinc 
dust  or  scrap  iron  with  slaked  lime,  hydrogen  is  liberated, 
and  an  oxide  of  the  metal  used  is  simultaneously  formed. 
This  method  is  frequently  used  for  preparing  hydrogen  in 
large  quantities  for  industrial  purposes. 

By  far  the  commonest  way  of  preparing  hydrogen  in  the 
laboratory  is  by  treating  zinc  with  dilute  sulphuric  acid. 
The  apparatus  used  for  this  purpose  is  shown  in  Fig.  5.  In 
this  reaction  there  is  formed,  besides  hydrogen,  a  white  salt 


Fia.  5. 


HYDROGEN 


17 


called  zinc  sulphate.  It  remains  in  solution,  and  may  be 
obtained  in  form  of  crystals  by  evaporating  the  solution  to 
a  small  bulk  and  allowing  it  to  cool.  We  may  express  the 
change  thus : — 

Sulphuric  Acid  +  Zinc  =  Hydrogen  +  Zinc  Sulphate. 

Instead  of  sulphuric  acid,  dilute  hydrochloric  acid  or  acetic 
acid  may  be  used.  Furthermore,  iron  may  be  substituted  for 
the  zinc,  in  which  case  hydrogen  and  corresponding  salts  of 
iron  are  formed.  Hydrogen  thus  obtained  is  never  quite  pure. 
The  impurities  present  in  ordinary  zinc  and  iron,  such  as  carbon, 
arsenic,  sulphur,  and  phosphorus,  combine  with  some  of  the 
hydrogen,  and  the  relatively  small  amounts  of  the  resulting 
gases  contaminate  the  larger  portion  of  hydrogen  which  remains. 
These  impurities  may  be  removed  by  passing  the  gas  through 
appropriate  absorbents.  Ordinary  cast  iron  usually  contains 
so  much  of  the  impurities  mentioned,  notably  of  carbon,  that, 
when  treated  with  an  acid,  the  hydrogen  liberated  is  con- 
taminated sufficiently  to  have  a  very  disagreeable  odor. 

Properties.  —  Hydrogen  is  the  lightest  of  all  known  sub- 
stances. It  is  a  colorless,  odorless,  tasteless  gas.  At  0°  and 
760  mm.  barometric 
pressure,  namely,  under 
standard  conditions, 
one  liter  weighs  0.08987 
gram.  It  is  14.388 
times  lighter  than  the 
air ;  in  other  words,  its 
specific  gravity  with 
respect  to  air  is  0.0695. 
Because  of  the  light- 
ness of  hydrogen,  jars 
containing  the  gas  are 
held  bottom  upward. 
Figure  6  shows  how  hydrogen  may  be  transferred  from  one  jar 
A  into  another  jar  B.  At  and  below  —  241°,  its  critical  tem- 
perature, hydrogen  may  be  liquefied  by  subjecting  it  to  pres- 
sure. At  —  241°  a  pressure  of  20  atmospheres  will  liquefy  the 
gas;  but  at  —252.5°  the  vapor  tension  of  liquid  hydrogen  is 
practically  one  atmosphere ;  that  is  to  say,  the  liquid  boils  at 


FIG. 


18 


OUTLINES   OF   CHEMISTRY 


the  last-named  temperature.  Liquid  hydrogen  is  clear  and 
colorless,  like  water,  and  has  a  specific  gravity  of  about  0.07. 
Solid  hydrogen  may  be  obtained  by  evaporating  liquid  hydrogen 
in  a  partial  vacuum.  The  melting  point  of  the  solid,  which 

consists  of  white  crystals, 
is  —  259°.  Its  power  to 
refract  light  is  6.5  times 
greater  than  that  of  air. 

On  account  of  its  light- 
ness, hydrogen  diffuses 
very  rapidly,  and  readily 
passes  through  porous  sub- 
stances like  unglazed  por- 
celain, brick,  mortar,  and 
paper.  The  rate  of  diffu- 
sion of  gases  is  inversely 
proportional  to  the  square 
roots  of  'their  densities ; 
hence,  air  diffuses  only 
V0.0695  or  0.2636  time 
as  fast  as  hydrogen.  The 
rapid  diffusion  of  hydrogen 
may  be  demonstrated  by 
means  of  the  apparatus 
shown  in  Fig.  7.  When 
the  unglazed  porcelain  cup 
A,  which  contains  air,  is 
surrounded  with  hydrogen 
gas,  which  is  passed  into 
the  inverted  vessel  B  by 
means  of  the  tube  C,  the 
hydrogen  diffuses  into  the 
porous  cup  A  much  faster 
than  the  air  diffuses  out. 
A  pressure  is  consequently 
produced  in  A  which  is  connected  with  the  Wolf  bottle  D 
containing  water;  and  this  pressure  forces  the  water  out  of 
the  tube  E  in  form  of  a  fountain. 

Hydrogen  is  but  slightly  soluble  in  water,  for  only  19  vol- 
umes are  absorbed  by  1000  volumes  of  water  at  15°.     Certain 


FIG.  7, 


HYDROGEN 


19 


solids  absorb  hydrogen  in  notable  quantities.  Freshly  ignited 
charcoal  absorbs  about  twice  its  volume  of  hydrogen.  Palla- 
dium absorbs  500  volumes,  platinum  49  volumes,  iron  19  vol- 
umes, gold  46  volumes,  copper  4.5  volumes,  nickel  17  volumes, 
aluminum  2.7  volumes,  lead  0.15  volume.  At  red  heat,  pal- 
ladium may  even  absorb  as  much  as  900  volumes,  according  to 
Graham.  This  power  of  solids  to  absorb  gases  is  sometimes 
termed  adsorption,  or  occlusion.  The  amount  absorbed  depends 
upon  the  specific  nature  of  the  solid  and  also  of  the  gas ;  and 
as  the  absorption  is  accompanied  with  changes  of  temperature 
and  of  volume,  it  is  clear  that  the  phenomenon  is  akin  to  the 
process  of  solution.  Furthermore,  hydrogen  passes  through 
iron  and  platinum  tubes  when  these  are  hot.  This  is  readily 
explained  by  the  fact  that  these  metals  absorb  the  gas. 

The  most  notable  characteristic  of  hydrogen  is  its  inflamma- 
bility. It  burns  readily  in  the  air  or  in  oxygen,  and  the  product 
formed  is  water.  This  can  easily  be  shown  by  holding  a  cold 
bell  jar  over  a  burning  jet  of  hydrogen  (Fig.  8).  The  water 


FIG.  8. 


formed  condenses  in  drops  on  the  sides  of  the  jar.  At  ordinary 
temperatures,  hydrogen  and  oxygen  do  not  act  on  each  other 
appreciably,  but  the  action  takes  place  when  the  gases  are 
heated  to  the  kindling  temperature,  which  is  about  615°,  ac- 
cording to  V.  Meyer.  The  hydrogen  flame  is  colorless.  To 
show  this  the  gas  must  be  burnt  from  a  platinum  jet,  which 
is  not  affected  during  the  process,  so  that  particles  of  foreign 
matter  do  not  get  into  the  flame  and  color  it.  When  a  glass 


20 


OUTLINES   OF  CHEMISTRY 


tube  is  held  over  a  hydrogen  flame,  as  shown  in  Fig.  9,  the 

column  of  air  in  the  tube  is  set  in  vibration,  thus  producing 

the  phenomenon  known  as  the  singing 
flame. 

The  hydrogen  flame  is  very  hot, 
which  is  evident  from  the  fact  that 
platinum,  which  fuses  above  1700°, 
will  melt  in  it.  The  burning  of  one 
gram  of  hydrogen  develops  about  34.5 
large  calories  of  heat,  which  is  enough 
heat  to  raise  the  temperature  of  345 
grams  of  water  from  0°  to  the  boiling 
point  or  to  melt  431  grams  of  ice. 

Hydrogen  is  not  poisonous,  but 
animals  would  suffocate  in  the  gas 
for  lack  of  oxygen,  without  which 
they  cannot  live.  Hydrogen  will  also 
not  support  ordinary  combustion. 
Thus,  when  a  lighted  candle  is  thrust 
into  a  jar  of  hydrogen  (Fig.  10),  the  gas 
at  the  mouth  of  the  jar  --- — -^ 

is  set  on  fire,  but  the 
flame   of    the   candle 
is  extinguished. 
At    room    temperatures    and    atmospheric 

pressure  hydrogen  is  rather  inert  chemically, 

combining  with  vigor  with  but  one  element ; 

namely,  fluorine.     But  in  the  sunlight,  hydro- 
gen   becomes    more    active,   notably    toward 

chlorine,    with    which    it    combines    readily, 

forming   hydrochloric  acid.     With    nitrogen, 

hydrogen    forms    ammonia ;     with    sulphur, 

hydrogen  sulphide;    with  carbon,  marsh  gas 

and    other    hydrocarbons.      A   compound   of 

hydrogen  with   another   element   is   called   a 

hydride.      Thus,    water    may    be    termed    a 

hydride  of  oxygen;  it  may  also  be  called  an 

oxide  of  hydrogen. 

Since  hydrogen  readily  unites  with  oxygen  at  elevated  tem- 
peratures, it  may  be  used  to  deprive  some  compounds  of  their 


FIG. 


FIG.  10. 


HYDROGEN 


oxygen  content.  So  when  copper  is  heated  in  the  air  it  turns 
black  because  of  union  of  the  surface  layers  with  oxygen  of  the 
air,  and  when  this  hot,  black  copper  oxide  is  now  brought  into  an 
atmosphere  of  hydrogen,  the  latter  gas  unites  with  the  oxygen, 
forming  water  and  bright,  metallic  copper.  The  process  of 
union  of  the  copper  with  oxygen  is  termed  oxidation,  whereas 
the  process  of  depriving  the  copper  oxide  of  its  oxygen  content 
is  called  reduction.  Oxidation  and 
reduction  are  thus  opposite  processes. 
At  higher  temperatures,  hydrogen  is 
similarly  able  to  reduce  quite  a 
number  of  oxides,  such  as  the  oxides 
of  lead,  iron,  mercury,  zinc,  and 
nickel.  On  account  of  its  power  to 
abstract  oxygen  from  compounds, 
hydrogen  is  called  a  reducing  agent. 
The  process  of  oxidation  and  reduc- 
tion of  copper  is  readily  illustrated 
by  means  of  the  apparatus  shown  in 
Figs.  11  and  12.  The  bright  copper 


FIG.  11. 


FIG.  12. 


crucible  is  heated  by  means  of  the  burner  as  shown  in  Fig.  11. 
The  black  oxide  of  copper  is  thus  formed  on  the  surface  of  the 
crucible.  The  flame  is  now  extinguished,  and  while  the 
crucible  is  still  quite  hot  it  is  brought  into  an  atmosphere  of 
hydrogen.  This  is  accomplished  by  holding  the  large  funnel, 
through  which  a  strong  current  of  hydrogen  is  being  passed, 
down  over  the  crucible  so  as  to  envelop  it  (Fig.  12);  the 
black  crucible  thus  quickly  assumes  a  bright  copper  color. 


22 


OUTLINES  OF  CHEMISTRY 


Uses.  —  Besides    being   employed   as   a   reducing   agent   in 

various  chemical  operations  in  the  laboratory  and  the  indus- 
tries, hydrogen  finds  use  on  account  of  its  light- 
ness  and  combustibility.  Its  lightness  makes  it 
specially  suitable  for  filling  balloons.  Hydrogen 
prepared  by  electrolysis  and  placed,  under  pres- 
sures of  100  to  150  atmospheres,  in  steel  cylinders 
(Fig.  13),  is  now  an  article  of  commerce.  Mixed 
with  carbon  monoxide,  hydrogen  forms  an  im- 
portant part  of  water  gas,  which  is  used  for 
heating  purposes.  Hydrogen  is  too  expensive 
for  use  in  ordinary  heating.  It  is  at  times 
employed  as  fuel  in  operations  where  very  high 
temperatures  are  required,  as  in  working  platinum 
and  other  metals  having  a  high  melting  point. 

Hydrogen  Equivalents  of  Metals.  — Whether  a 
metal  liberates  hydrogen  from 
water  itself  or  from  various 
dilute  acids,  the  amount  of  hydro- 
gen which  a  given  weight  of  a 
certain  metal  will  set  free  is  the 
same.  The  amount  of  hydrogen 
which  a  definite  weight  of  a 

given  metal  is  able  to  liberate  may  be  ascer- 
tained by  means  of  the  apparatus  shown  in 

Fig.  14.     Water  is  first  placed  in  the  beaker  B. 

The  graduated  tube  A  is  placed  into  the  water 

as  shown,  and   by  means  of   suction  at   the 

upper  end  of  A,  while  the  cock  0  is  open,  the 

water  is  drawn  up  the  tube  till  it  is  filled  to 

a  little  above  the  cock,  which  is  then  closed. 

The  tube  A  is  then  raised  slightly  and  its 

lower  end  placed   into  the  little  crucible  D, 

which  contains  a  weighed  quantity  of  a  metal, 

say  magnesium.     The  upper  "end  of  A  is  then 

filled   with   dilute   acid,    which   by   carefully 

opening  the  cock  0  is  allowed  to  flow  down 

upon  the  metal.     The  cock  0  is  closed  before 

all  the  acid  has  passed  the  cock  so  as  to  avoid  admitting  air 

into  the  graduated  tube.     After  the  acid  has  dissolved  the 


o-c 


FIG.  13. 


FIG.  14. 


HYDROGEN  23 

metal  completely,  the  level  of  the  liquid  in  the  tube  A 
is  adjusted  so  that  it  is  the  same  as  that  in  the  beaker. 
The  volume  of  the  hydrogen  is  noted,  the  temperature  and 
barometric  pressure  are  taken,  and  from  these  data  the  volume 
of  the  hydrogen  under  standard  conditions  is  computed. 
Knowing  this  and  the  weight  of  one  liter  of  hydrogen,  the 
weight  of  the  hydrogen  liberated  may  readily  be  calculated. 
The  result  would  be  the  weight  of  hydrogen  displaced  from 
the  acid  by  the  given  weight  of  magnesium,  and  from  this  the 
amount  of  magnesium  required  to  liberate  1  gram  of  hydro- 
gen can  easily  be  found.  An  experiment  of  this  kind  yields 
the  result  that  it  requires  12.16  grams  of  magnesium  to  liberate 
1  gram  of  hydrogen.  Similarly,  it  has  been  found  that  23.00 
grams  of  sodium,  or  39.10  grams  of  potassium,  or  9.03  grams 
of  aluminum,  or  27.9  grams  of  iron,  or  59.5  grams  of  tin,  or  32.7 
grams  of  zinc  are  required  to  set  free  1  gram  of  hydrogen.  The 
quantities  mentioned  are  called  the  hydrogen  equivalents  of  the  re- 
spective metals ;  or  sometimes  they  are  simply  spoken  of  as  the 
chemical  equivalents.  It  is  evident  that  the  amounts  of  the  vari- 
ous metals  that  are  chemically  equivalent  to  1  gram  of  hydrogen 
are  very  different.  The  chemical  equivalents  of  the  elements 
are  of  great  importance,  and  they  will  be  referred  to  again  later. 
When  each  of  the  metals  above  mentioned  acts  upon  dilute 
hydrochloric  acid,  it  is  evident,  from  even  a  rough  observation, 
that  the  rate  with  which  the  different  metals  liberate  hydrogen 
varies  greatly.  Arranging  these  metals  in  the  order  of  rapidity 
with  which  they  react  with  dilute  acid,  wTe  have  :  potassium, 
sodium,  magnesium,  aluminum,  zinc,  iron,  and  tin,  the  action 
being  strongest  in  the  case  of  potassium,  and  weakest  in  the  case 
of  tin.  This  gives  us  an  idea  of  the  relative  affinity  or  chemi- 
cal attraction  that  exists  between  these  metals  and  the  dilute 
aqueous  solution  of  the  acid  used,  or  rather  between  the  metals 
and  that  part  of  the  aqueous  acids  with  which  the  displaced 
hydrogen  was  combined.  By  measuring  accurately,  at  con- 
stant temperature,  the  rate  of  the  liberation  of  hydrogen  per 
minute  when  one  and  the  same  area  of  the  different  metals  acts 
on  samples  of  the  same  dilute  acid  solution,  the  relative  affini- 
ties of  the  metals  for  the  acid  may  be  determined  ;  for  the  rate 
with  which  a  chemical  reaction  proceeds  is  proportional  to  the 
chemical  affinity  that  comes  into  play. 


CHAPTER   III 

OXYGEN 

History. —  Oxygen  was  discovered  in  1774  by  Joseph  Priestley, 
who  liberated  the  gas  by  heating  the  red  oxide  of  mercury.  It 
was  independently  discovered  in  1773  by  Scheele,  but  he  did 
not  publish  his  work  till  1775.  Lavoisier,  who  found  the  dis- 
covery of  the  gas  of  particular  interest  in  connection  with  his 
studies  of  the  process  of  combustion,  named  the  element  oxygen, 
from  the  Greek  words  meaning  acid  and  to  generate.  He  found 
that  the  union  of  oxygen  with  such  elements  as  sulphur,  nitrogen, 
and  arsenic  produced  substances  that  were  sour  to  the  taste, 
and  in  general  behaved  like  other  well-known  acids.  His  con- 
clusion was  that  oxygen  is  an  essential  constituent  of  all  acids, 
but  later  work  has  shown  this  to  be  erroneous. 

Occurrence. —  Oxygen  is  the  most  abundant  element  on  the 
earth.  The  atmosphere  contains  about  21  per  cent  of  free  oxy- 
gen by  volume.  Water  contains  88.88  per  cent  of  oxygen  by 
weight,  and  the  rocks  of  the  earth's  crust  contain  from  44  to 
48  per  cent.  It  is  present  in  all  animals  and  plants,  in  which  it 
occurs  in  combination  with  hydrogen  and  carbon,  and  also  with 
hydrogen,  carbon,  and  nitrogen. 

Preparation. —  (1)  When  liquid  air  is  allowed  to  evaporate, 
the  nitrogen,  which  is  more  volatile  than  the  oxygen,  passes  off 
first,  and  thus  a  considerable  portion  of  the  oxygen  is  left  in  the 
container,  approximately  free  from  nitrogen.  (2)  By  electroly- 
sis of  water,  acidified  with  sulphuric  acid,  two  volumes  of  hydro- 
gen and  one  volume  of  oxygen  are  produced.  (3)  By  heating 
red  oxide  of  mercury,  this  compound  is  decomposed,  yielding 
oxygen  and  mercury;  similarly,  the  oxide  of  silver  may  be 
decomposed  by  heat.  Again,  the  peroxides  of  manganese,  lead, 
and  barium  give  off  a  portion  of  their  oxygen  on  heating  them. 
The  peroxides  of  these  metals  also  evolve  oxygen  when  heated 
with  sulphuric  acid.  (4)  Certain  salts  rich  in  oxygen  give  off 
their  oxygen  content  either  in  part  or  entirely  upon  being  heated. 

24 


OXYGEN  25 

Thus,  saltpeter  yields  oxygen  and  potassium  nitrite  on  ignition, 
and  potassium  chlorate  when  heated  yields  oxygen  and  potassium 
chloride.  The  latter  method  is  very  commonly  used  for  preparing 
oxygen  for  laboratory  purposes.  One  hundred  grams  of  potassium 
chlorate  yield  about  39  grams  of  oxygen.  In  the  process  of 
heating  potassium  chlorate,  potassium  perchlorate  first  forms, 
and  this  upon  further  heating  breaks  down  into  oxygen  and 
potassium  chloride.  (5)  By  treating  a  solution  of  hydrogen 
peroxide,  acidified  with  sulphuric  acid,  with  potassium  perman- 
ganate or  potassium  bichromate,  oxygen  is  evolved.  This 
method  is  very  convenient  for  laboratory  purposes.  (6)  When 
bleaching  powder  acts  on  peroxide  of  hydrogen,  oxygen  is 
evolved.  (7)  Barium  oxide  when  heated  in  the  air  to  about 
500°  takes  on  oxygen,  forming  barium  peroxide.  The  latter  on 
being  heated  up  to  1000°  parts  with  half  of  its  oxj^gen,  forming 
the  original  barium  oxide,  and  the  process  can  then  be  repeated. 
This  is  known  as  Brin's  process.  It  will  be  seen  that  it  is  a  con- 
venient method  of  preparing  oxygen  from  the  air.  It  is  used 
for  preparing  oxygen  for  commercial  purposes.  (8)  The  green 
leaves  of  plants  in  the  sunlight  decompose  carbon  dioxide  and 
water,  forming  starch  and  oxygen.  Large  quantities  of  oxygen 
are  thus  supplied  to  the  atmosphere. 

Properties.  —  Oxygen  is  a  colorless,  odorless,  tasteless  gas. 
It  is  1.10  times  as  heavy  as  air.  One  liter  under  standard 
conditions  (0°  and  760  mm.)  weighs  1.4290  grams.  Its  power 
to  refract  light  is  only  0.8616  time  that  of  air.  The  gas  may 
be  liquefied  at  and  below  — 119°,  its  critical  temperature.  At 
—  119°  a  pressure  of  fifty  atmospheres  is  required  to  liquefy 
oxygen.  This  pressure  is  consequently  the  critical  pressure. 
Liquid  oxygen  is  a  light  blue,  mobile  liquid  which  boils  at 
— 182.5°  under  atmospheric  pressure.  It  is  attracted  by  a  mag- 
net. At  -182.5°  the  specific  gravity  of  the  liquid  is  1.1315. 
By  means  of  liquid  hydrogen,  Dewar  froze  oxygen  to  a  pale 
blue,  snowlike  solid,  whose  melting  point  is  —  227°. 

Oxygen  is  slightly  soluble  in  water.  At  0°  and  atmospheric 
pressure  100  volumes  of  water  dissolve  four  volumes  of  oxygen, 
while  at  15°,  3.4  volumes  of  the  gas  are  absorbed.  Oxygen 
may  consequently  be  collected  over  water. 

Chemically,  oxygen  is  a  very  active  substance  combining 
directly  with  all  known  elements,  the  only  exceptions  being 


26 


OUTLINES  OF  CHEMISTRY 


FIG.  15. 


fluorine  and  the  gases  of  the  argon  group,  namely,  helium, 
neon,  argon,  krypton,  and  xenon.  The  compounds  of  the 
elements  with  oxygen  are  called  oxides.  At  ordinary  tempera- 
tures, oxygen  unites  but  slowly  with  most  substances.  Thus, 
the  rusting  of  iron  consists  of  a  slow  union 
with  oxygen  of  the  air.  Sodium  is  oxi- 
dized quite  rapidly  on  exposure  to  air  or 
oxygen  at  room  temperature,  while  in  the 
case  of  wood,  charcoal,  or  sulphur,  the 
union  with  oxygen  at  ordinary  tempera- 
tures proceeds  very  slowly  indeed.  How- 
ever, at  elevated  temperatures  all  of  these 
substances  combine  readily  and  vigorously 
with  oxygen,  with  concomitant  evolution 
of  heat  and  light.  This  process  is  termed 
combustion.  All  chemical  processes  which 
proceed  with  the  evolution  of  light  and 
heat  may,  in  general,  be  called  cases  of 
combustion  ;  ordinarily,  however,  the 

term   is   applied   to   union  with   oxygen.     In   an   atmosphere 
of  the  latter   gas,  iron  will  burn  with  brilliant  scintillations 
(Fig.  15)  and  evolution  of  much  heat.     The  product  formed  is 
an  oxide  of  iron  of  a  reddish  brown  color. 
Phosphorus   burns  brilliantly  in  oxygen, 
forming   phosphoric   oxide,   consisting   of 
white  fumes  which  condense  on  the  sides 
of    the    container.       On   moistening   this 
white  solid  with  water,  a  solution  of  phos- 
phoric acid  is  formed.     This  solution  is 
sour  and  turns  blue  litmus  red.     Carbon 
burns  in  oxygen  to  carbon  dioxide ;   sul- 
phur burns  to  sulphur  dioxide  (Fig.  16). 
These  gases,  too,  form  acids  when  treated 
with  water.     The  oxides  of   phosphorus, 
carbon,    and    sulphur    are    consequently 
acid -forming     oxides.       They     are.    also 
spoken  of  as  acid  anhydrides;  that  is,  the  acids  minus  water. 
Sodium  when  burned  in  oxygen  forms  a  white  powder,  called 
sodium  oxide,  which  readily  dissolves  in  water,  yielding  a  solu- 
tion which  is  alkaline  to  the  taste,  turns  red  litmus  blue,  and 


-—    -?~V:^* 


FIG.  16. 


OXYGEN  27 

feels  slippery  to  the  fingers.  It  is  an  alkali  or  base,  and  is 
capable  of  reacting  with  acids,  forming  salts  whose  aqueous 
solutions  have  no  effect  on  litmus,  i.e.  they  are  neutral. 
Potassium  and  calcium  also  burn  readily  in  oxygen,  forming 
the  oxides  of  potassium  and  calcium.  These  are  white  caustic 
substances  which  resemble  the  oxide  of'  sodium.  The  oxide  of 
calcium  is  ordinary  lime.  The  oxides  of  potassium  and  calcium 
are  caustic  alkalies.  Other  oxides,  like  those  of  zinc,  iron,  and 
lead,  are  insoluble  in  water.  They  are  consequently  tasteless 
and  do  not  affect  litmus.  The  oxides  of  most  metals  can  be 
formed  by  direct  union  with  oxygen.  Some  metals,  like  gold 
and  platinum,  do  not  burn  in  oxygen,  but  their  oxides  may  be 
formed  indirectly  by  double  decomposition.  On  heating  such 
oxides,  they  yield  the  metal  and  oxygen. 

Combustion  in  the  Air.  —  The  combustion  of  substances  in 
the  air  yields  precisely  the  same  products  as  combustion  in  oxy- 
gen. Indeed,  the  process  of  burning  substances  in  the  air  is  in 
all  respects,  except  in  brilliancy,  rapidity,  and  vigor,  like  that 
of  burning  them  in  oxygen.  As  the  oxygen  of  the  air  is  di- 
luted with  four  times  its  volume  of  nitrogen,  which  latter  gas 
is  rather  inert  in  character,  it  is  quite  natural  that  combustion 
in  the  air  should  go  on  less  vigorously  than  in  oxygen.  The 
total  energy  liberated  as  heat  is  the  same,  however,  whether  the 
oxidation  of  a  substance  takes  place  rapidly  in  pure  oxygen,  or 
less  rapidly  in  the  air,  or  extremely  slowly  at  ordinary  tempera- 
tures in  the  air. 

Kindling  Temperature  and  Temperature  of  Combustion.  —  In 
order  to  burn  a  substance  in  oxygen,  it  must  be  heated  to  a 
certain  minimum  temperature  at  which  it  will  burst  into  flame. 
This  temperature,  which  is  very  different  for  different  sub- 
stances, is  called  the  kindling  temperature.  Thus,  phosphorus 
catches  fire  at  a  much  lower  temperature  than  sulphur,  and  the 
latter  ignites  at  a  lower  temperature  than  wood. 

The  highest  temperature  attainable  during  the  process  of 
combustion  of  a  substance  is  sometimes  called  the  temperature 
of  combustion.  It  varies  greatly  with  the  nature  of  the  sub- 
stance. It  is  higher  in  pure  oxygen  than  in  air,  and  higher  in 
compressed  oxygen  than  in  that  gas  at  atmospheric  pressure. 
The  temperature  of  combustion  is  generally  very  much  higher 
than  the  kindling  temperature. 


28  OUTLINES   OF  CHEMISTRY 

Heat  of  Combustion.  —  The  heat  evolved  during  the  combus- 
tion of  a  substance  is  called  its  heat  of  combustion.  As  above 
stated,  it  is  the  same  whether  the  combustion  goes  on  rapidly 
or  slowly,  though  the  maximum  temperature  reached  during  the 
process  of  combustion  varies  greatly  under  different  conditions. 

The  unit  of  heat  is  the  calorie.  The  small  calorie  is  the 
amount  of  heat  required  to  raise  1  gram  of  water  1  degree ; 
the  large  calorie  is  1000  times  as  large,  i.e.  it  is  the  amount  of 
heat  required  to  raise  1000  grams  of  water  1  degree  in  tern 
perature.  It  is  very  important  to  ascertain  the  heat  of  com- 
bustion of  various  substances,  not  only  for  purely  scientific 
purposes,  but  also  for  the  determination  of  the  relative  value  of 
fuels  and  certain  classes  of  food.  Heats  of  combustion  will 
consequently  receive  special  consideration  in  the  chapter  on 
thermochemistry. 

Different  Stages  of  Oxidation.  —  While  it  is  true  that  combus- 
tion in  the  air  or  in  oxygen  is  essentially  the  same  process, 
except  as  to  rapidity,  it  not  infrequently  happens  that  when  a 
substance  is  burnt  in  an  excess  of  oxygen,  more  of  the  latter 
enters  into  the  oxides  formed  than  when  the  burning  proceeds 
in  the  air.  Thus,  when  iron  is  oxidized  by  heating  it  in  the 
air,  a  black  oxide  is  formed  which  is  magnetic  in  character,  and 
which  consists  of  72.38  per  cent  iron  and  27.62  per  cent  oxy- 
gen ;  whereas  when  iron  is  burned  in  oxygen,  there  is  formed 
mainly  a  reddish  brown  oxide  of  iron  which  is  practically  non- 
magnetic, and  which  contains  69.96  per  cent  iron  and  30.04  per 
cent  oxygen.  By  carefully  heating  the  latter  oxide  in  a  cur- 
rent of  hydrogen  at  500°  a  black  oxide  may  be  obtained  which 
consists  of  77.75  per  cent  iron  and  22.25  per  cent  oxygen. 
Writing  the  composition  of  these  oxides  of  iron,  the  only  ones 
known,  in  form  of  a  table,  we  have  as  follows :  — 


PEE  CENT  IRON 

PER  CENT  OXYGEN 

PARTS  OXYGEN  TO  77.75  PARTS  IRON 

(1)    72.38 

(2)  69.96 
(3)  77.75 

27.62 
30.04 
22.25 

29.67 
33.38 
22.25 

In  the  third  column  are  placed  the  amounts  of  oxygen  by 
weight  combined  with  one  and  the  same  amount  of  iron;  namely, 


OXYGEN 


29 


77.75  parts.  The  latter  figure  was  chosen  simply  for  conven- 
ience, as  it  represents  the  percentage  of  iron  in  the  oxide  poor- 
est in  oxygen.  Now,  inspecting  the  table,  we  see  that 

29.67:22.25  =  4:3, 

and  that  33.38:22.25  =  3:2. 

This  means  that  in  these  three  different  oxides  of  iron  the 
amounts  of  oxygen  that  are  combined  with  one  and  the  same 
amount  of  iron  are  simple,  rational  multiples  of  one  another. 
This  being  the  case,  had  we  calculated  the  amounts  of  iron 
combined  in  these  oxides  with  one  and  the  same  amount  of 
oxygen,  we  should  have  found  that  these  amounts  of  iron  are 
also  simple,  rational  multiples  of  one  another. 

Again,  there  are  five  different  oxides  of  lead  known.  These 
are  as  follows:  (1)  lead  suboxide,  a  black  substance  formed 
when  lead  is  heated  at  its  melting  point  in  the  air ;  (2)  lith- 
arge, a  yellow  powder  formed  when  lead  is  very  strongly  heated 
in  air ;  (3)  lead  sesquioxide,  an  orange-yellow  powder  formed 
when  bleaching  powder  acts  on  litharge  dissolved  in  caustic 
potash ;  (4)  red  lead  or  minium,  a  bright  red  powder,  which 
may  be  obtained  by  heating  litharge  in  the  air  at  a  temperature 
not  above  450°;  and  (5)  lead  peroxide,  a  brown  powder,  which 
may  be  prepared  by  treating  red  lead  with  dilute  nitric  acid. 
The  percentage  composition  of  these  oxides  is  as  follows  :  — 


NAME 

PER  CENT  LEAD 

PER  CENT  OXYGEN 

PARTS  LEAD  TO  3.72 
PARTS  OXYGEN 

(1)  Lead  suboxide  .     .     .' 
(2)  Litharge    . 

96.28 
92  82 

3.72 

7.18 

96.28 
48.14 

(3)  Lead  sesquioxide    .     . 
(4)  Red  lead    .     . 

89.61 
90  65 

10.39 
9.35 

32.09 
36.11 

(5)  Lead  peroxide    .     .     . 

86.60 

13.40 

24.07 

In  the  last  column  we  have  the  amount  of  lead  combined 
with  3.72  parts  of  oxygen  in  each  of  the  oxides.  Comparing 
the  figures  in  the  last  column  we  note  as  follows  :  — 


(1)  and  (2) 
(1)  and  (3) 
(1)  and  (4) 
(1)  and  (5) 


96.28:48.14  =  2:1, 
96.28:32.09  =  3:1, 
96.28:36.11  =  8:3, 
96.28:24.07  =  4:1. 


30  OUTLINES   OF  CHEMISTRY 

Thus  we  see  that  in  the  five  oxides  of  lead,  the  amounts  of 
lead  combined  with  one  and  the  same  amount  of  oxygen  are 
simple  multiples  of  one  another.  Obviously  the  amounts  of 
oxygen  which  in  these  oxides  are  combined  with  one  and  the 
same  amount  of  lead  are  also  simple  multiples  of  one  another. 

Law  of  Multiple  Proportions.  —  These  results  of  the  quantita- 
tive study  of  the  composition  of  the  oxides  of  iron  and  lead  are 
typical  of  a  large  number  of  similar  cases.  It  has  been  found 
to  be  general,  that  whenever  two  elements  form  more  than  one 
compound  with  each  other,  the  amounts  by  weight  of  the  one  that 
are  united  with  one  and  the  same  weight  of  the  other  are  simple 
rational  multiples  of  one  another.  This  is  the  law  of  multiple 
proportions.  It  was  discovered  by  John  Dalton  about  1806. 
Many  careful  analyses  of  various  compounds  have  since  yielded 
results  confirming  this  law,  which  is  of  fundamental  importance 
in  chemistry.  As  we  proceed,  we  shall  meet  numerous  addi- 
tional instances  illustrating  the  law  of  multiple  proportions. 

Role  of  Oxygen  in  Respiration.  —  Oxygen  is  necessary  for  all 
animal  life.  If  the  oxygen  supply  is  cut  off  from  an  animal,  it 
soon  dies  from  suffocation.  Pure  oxygen  may  be  inhaled  with- 
out evil  effects  for  a  while.  An  animal  placed  in  oxygen 
shows  invigoration  by  its  more  lively  movements;  but  after  a 
while  febrile  symptoms  appear,  and  a  reaction  sets  in  which  may 
cause  death.  The  air  as  it  enters  the  lungs  is  virtually  oxygen 
diluted  with  four  times  its  volume  of  nitrogen.  It  is  the 
oxygen  only  that  is  absorbed  by  the  membranes  of  the  lungs. 
Furthermore,  only  4  to  5  per  cent  of  the  oxygen  contained  in 
the  air  is  thus  absorbed  in  the  process  of  respiration.  The  ex- 
haled air  contains  water  and  also  about  3  to  4  per  cent  of  car- 
bon dioxide,  gained  from  the  body. 

The  oxygen  from  the  air  passes  through  the  membranes  of 
the  lungs,  into  the  blood,  where  it  is  taken  up  by  the  blood 
corpuscles.  The  latter  contain  hemoglobin,  a  crystalline  sub- 
stance which  unites  with  oxygen,  forming  oxyhemoglobin, 
which  has  a  red  color,  giving  a  bright  appearance  to  arterial 
blood.  As  oxyhemoglobin  attached  to  the  blood  corpuscles, 
the  circulation  carries  the  oxygen  to  all  parts  of  the  body, 
where  it  is  given  off,  entering  into  various  combinations  with 
the  tissues.  As  the  blood  is  thus  deprived  of  oxygen,  carbon 
dioxide,  which  is  formed  during  the  oxidation  of  the  tissues,  is 


OXYGEN 


31 


taken  up  and  carried  to  the  lungs,  where  it  is  exhaled  and  ex- 
changed for  oxygen.  The  blood  deprived  of  a  portion  of  its 
oxygen  and  laden  with  carbon  dioxide  is  so-called  venous  blood. 
It  is  dark  in  color  instead  of  bright  red.  On  discharging  its 
carbon  dioxide  and  taking  on  oxygen,  it  is  converted  into 
so-called  arterial  blood,  which  is  bright  red.  All  of  these  pro- 
cesses go  on  much  more  rapidly  and  vigorously  in  an  atmos- 
phere of  pure  oxygen  than  in  air.  It  is  for  this  reason  that 
animals  succumb  in  oxygen  ;  they  are  destroyed  by  the  too 
rapid  changes.  On  the  other  hand,  if  the  supply  of  oxygen  is 
unduly  diminished,  the  transformations  described,  which  are 
necessary  for  life,  cannot  go  on  and  the  animal  dies  of  suffoca- 
tion. As  stated  above,  pure  oxygen  may  be  breathed  for  a  time; 
it  is  frequently  administered  to  patients  who  are  suffering  de- 
pression because  of  difficulty  experienced  in  breathing. 

Fishes  derive  their  supply  of  oxygen  by  means  of  their  gills 
from  the  oxygen  dissolved  in  the  water. 

In  the  respiration  of  plants,  carbon  dioxide  is  taken  up  by 
the  green  leaves  in  the  sunlight,  and  oxygen  is  exhaled.  In  the 
leaf,  starch  is  simultaneously  formed,  as  carbon  dioxide  and 
water  act  on  each  other  with  elimination  of  oxygen.  Thus, 
while  animals  are  using  up  oxygen  in  breathing  and  are  giving 
off  carbon  dioxide,  plants  are  taking  up  the  latter  gas  and  re- 
turning oxygen  to  the  air. 

Oxyhydrogen  Blowpipe.  —  When  a  jet  of  hydrogen  is  burned 
in  the  air,  a  high  temperature  is  developed  ;  this  may  be 


FIG.  17. 


further  increased  by  burning  the  jet  in  oxygen,  or  by  supplying 
oxygen  to  the  jet  of  hydrogen  as  it  burns.  The  oxy hydrogen 
blowpipe  (Fig.  17)  is  an  arrangement  for  securing  very  high 
temperatures.  As  a  rule  the  burner  is  made  of  brass.  Hydro- 
gen passes  in  as  shown  and  issues  at  the  tip,  where  the  jet  is 


32 


OUTLINES  OF   CHEMISTRY 


lighted.  Oxygen  is  then  passed  in  as  indicated,  and  thus  the 
gases  do  not  mix  except  in  the  jet  itself.  In  this  way  explosions 
are  avoided.  The  oxyhydrogen  flame  readily  fuses  platinum 
or  silica,  and  is  used  in  working  such  refractory  materials. 
When  the  jet  is  directed  against  a  piece  of  lime,  the  latter  is 
heated  to  incandescence,  producing  a  very  intense  white  light, 
known  as  Drummond's  lime  light.  This  is  used  at  times  in 
projection  lanterns,  and  for  signaling  purposes  where  a  very 
intense  light  is  required. 

Detonating  Gas.  —  We  have  seen  that  when  water  is  decom- 
posed by  electrolysis,  two  volumes  of  hydrogen  and  one  volume 
of  oxygen  are  produced.  A  mixture  of  these  two  gases  in  the 
proportions  mentioned  is  highly  explosive  when  ignited,  for 
water  is  formed  which,  by  the  intense  heat  generated,  is  at  once 
converted  into  steam,  thus  producing  the  explosion.  The  ex- 
plosive character  of  oxyhydrogen  gas  may  be  demonstrated  in 
a  harmless  way  by  making  soap  bubbles  filled  with  the  gas  and 
then  igniting  them.  Not  too  large  a  quantity  of  the  gas  should 
be  exploded  at  once  in  a  room,  for  the  report  is 
very  loud  and  may  rupture  the  eardrum. 

Combustion  of  Oxygen  in  Hydrogen.  —  It  has 
been  mentioned  that  a  jet  of  hydrogen  will  burn 
in  an  atmosphere  of  oxygen,  developing  intense 
heat.  It  is  equally  possible  to  burn  a  jet  of 
oxygen  in  an  atmosphere  of  hydrogen  (Fig.  18). 
The  hydrogen  is  first  lighted  at  the  mouth  of  the 
cylinder,  and  a  jet  of  oxygen  is  then  introduced. 
It  ignites  and  continues  to  burn  in  the  atmosphere 
of  hydrogen  as  shown.  The  fact  that  either  of 
these  two  gases  may  be  burned  in  an  atmosphere 
of  the  other  shows  the  real  nature  of  combustion, 
which  consists  of  a  chemical  union  of  the  two 
gases.  The  product  formed  is,  of  course,  water  in 
either  case. 

Earlier  Views  of  Combustion.  —  That  the  combus- 
tion of  substances  in  the  air  is  a  process  of  oxida- 
tion was  not  recognized  till  Lavoisier  showed  it  to  be  true  by 
experiment.  Before  Lavoisier,  the  view  prevailed  that  when  a 
substance  is  burned  a  subtile  principle  flies  out  of  it.  This 
notion  dates  back  to  antiquity.  It  was  probably  suggested  by 


H 


OXYGEN  83 

the  rising  of  the  smoke  of  ordinary  fires.  It  was  Georg  Ernst 
Stahl  (1660-1734),  professor  of  medicine  at  the  University  of 
Halle,  who  first  formulated  a  definite  theory  of  combustion.  He 
called  the  subtile  principle,  which  he  assumed  flies  out  of  bodies 
on  burning  them,  phlogiston,  which  means  that  which  is  com- 
bustible. So,  for  instance,  when  mercury  is  heated  in  the  air 
to  500°  a  red  powder  results,  which,  according  to  Stahl's  view, 
would  be  dephlogisticated  mercury.  Similarly  he  looked  upon 
other  oxides  as  bodies  that  had  been  deprived  of  phlogiston. 
Anything  that  was  combustible  contained  phlogiston.  Thus 
carbon  was  considered  very  rich  in  phlogiston.  By  heating, 
for  example,  dephlogisticated  lead  (yellow  oxide  of  lead)  with 
carbon,  the  latter  would  give  off  phlogiston  to  the  yellow  pow- 
der and  thus  change  it  back  to  lead.  In  general,  what  we  now 
term  oxidation  was  regarded  as  dephlogistication,  and  what 
we  call  reduction  was  regarded  as  a  process  of  taking  on  phlo- 
giston. The  phlogistic  theory  dominated  chemistry  in  the 
eighteenth  century ;  and,  indeed,  many  chemical  changes,  and 
among  them  rather  complicated  ones,  could  in  a  way  be  ex- 
plained by  means  of  the  theory.  In  fact,  Cavendish,  Priestley, 
and  Scheele  adhered  to  the  phlogistic  theory. 

It  was  known  to  the  adherents  of  the  phlogistic  view  that 
when  metals  are  calcined  by  heating  them  in  the  air,  the  result- 
ing powder  is  heavier  than  the  original  metal.  In  fact,  this 
was  known  even  a  hundred  years  before  the  phlogistic  theory 
was  promulgated  ;  but  it  was  not  regarded  as  an  especially  vital 
fact  in  forming  a  correct  view  of  combustion.  It  was  not  an  age 
of  careful  quantitative  experimentation,  and  the  value  of  facts 
established  by  accurate  measurements  was  frequently  not  seri- 
ously considered.  And  so  it  was  that  when  Lavoisier  pointed 
out  that  metals  grow  heavier  when  burned  in  the  air,  and  argued 
that  this  means  that  something  is  added  to  the  metal  rather 
than  subtracted  from  it  during  the  process,  his  argument  did 
not  meet  with  favor,  even  on  the  part  of  the  discoverers  of  oxy- 
gen themselves.  The  adherents  of  the  phlogistic  view  argued 
that  the  fact  that  substances  increase  in  weight  when  burned 
could  not  serve  to  prove  that  something,  namely  phlogiston, 
might  not  also  fly  out  of  the  substances  during  the  process  of 
combustion.  In  order  to  explain  the  fact  that  substances  grow 
heavier  when  burned,  some  of  the  followers  of  Stahl  even 


34 


OUTLINES   OF   CHEMISTRY 


suggested  that  phlogiston  might  be  a  substance  of  negative 
weight. 

Antoine  Laurent  Lavoisier  (1743-1794),  the  founder  of  mod- 
ern chemistry,  laid  great  stress  upon  the  increase  in  weight  of 
substances  during  combustion,  and  when  oxygen  was  discovered 
by  Scheele  and  Priestley  he  actually  demonstrated  that  it  is 
this  gas  which  unites  with  bodies  when  they  are  burned  in 
the  air.  Thus,  he  heated  a  quantity  of  mercury  in  a  retort 
(Fig.  19)  in  contact  with  air  for  twelve  days.  The  end  of  the 


FIG.  19. 

retort  opened  into  a  bell  jar,  the  opening  of  which  was  shut  off 
from  the  outer  air  by  means  of  mercury,  as  shown.  The  total 
volume  of  the  air  in  the  retort  and  bell  jar  was  about  one  liter. 
After  the  apparatus  had  cooled,  it  was  found  that  a  diminution 
of  volume  of  the  air  in  the  apparatus  amounting  to  about  170  cc. 
had  taken  place.  From  the  calcined  mercury,  which  he  col- 
lected, he  obtained  160'  cc.  oxygen  by  heating ;  and  thus  he 
showed  by  synthesis  and  analysis  the  real  nature  of  calcined 
mercury.  He  further  demonstrated  that  carbon  unites  with 
the  oxygen  of  certain  metallic  oxides  when  heated,  and  that  thus 
the  metal  itself  is  prepared  by  subtraction  of  oxygen  from  the 
calcined  metal  rather  than  by  the  addition  of  phlogiston  to 
it.  The  views  of  Lavoisier  were  stoutly  opposed  by  the  follow- 


OXYGEN  35 

ers  of  the  phlogistic  theory.  However,  facts  began  to  increase 
in  favor  of  Lavoisier's  explanations,  and  when  Cavendish 
showed  that  water  is  formed  when  hydrogen  and  oxygen  unite 
chemically,  the  former's  views  soon  triumphed.  Whereas  the 
followers  of  phlogiston  regarded  the  metals  and  other  combusti- 
ble elements  as  compound  bodies  containing  phlogiston,  we  now 
look  upon  them  as  simple  bodies  capable  of  uniting  with  oxy- 
gen under  proper  conditions. 


CHAPTER  IV 

WATER 

Occurrence. —  Water  is  found  in  oceans,  lakes,  and  rivers,  in 
the  soil  and  in  the  atmosphere.  It  occurs  in  the  solid,  liquid, 
and  vapor  states.  As  snow  and  ice  it  covers  the  vast  fields  of 
the  polar  regions,  the  highest  mountain  peaks,  and,  during  the 
winter,  large  areas  of  the  temperate  zones.  Falling  in  form  of 
rain,  snow,  and  hail,  water  permeates  the  soil  and  forms  springs, 
lakes,  and  rivers  that  carry  it  to  the  sea.  In  the  atmosphere, 
it  exists  as  vapor  which  by  condensation  may  form  fogs  and 
clouds.  The  amount  of  aqueous  vapor  that  the  air  may  hold 
varies  with  the  temperature.  One  million  liters  of  air  satu- 
rated with  water  vapor  at  0°  contain  4800  grams  of  water, 
while  at  20°  and  at  30°  this  amount  of  air  will  take  up  17,000 
and  29,840  grams  of  water,  respectively.  Ordinarily,  air  is 
saturated  with  water  vapor  to  but  two  thirds  of  its  capacity. 
When  the  moisture  content  of  the  air  reaches  but  four  tenths 
of  its  capacity,  the  air  feels  dry,  whereas  it  requires  nearly 
double  this  amount  of  humidity  to  cause  the  sensation  of  damp- 
ness. In  all  plants  and  animals,  water  is  found  in  relatively 
large  quantities.  Usually  organisms  are  made  up  of  over  fifty 
per  cent  of  water.  Many  minerals,  salts,  and  manufactured 
products  contain  water  more  or  less  loosely  bound. 

Preparation.  —  Water  is  formed  not  only  when  hydrogen 
and  oxygen  gases  unite,  but  also  when  hydrogen  acts  on  vari- 
ous oxides  at  high  temperatures,  and  when  compounds  contain- 
ing hydrogen  are  oxidized.  It  forms  during  the  process  of  the 
oxidation  of  the  tissues  of  organic  beings,  and  together  with 
carbon  dioxide  is  exhaled  by  animals.  All  natural  waters  are, 
chemically  speaking,  impure.  Rain  water  is  the  purest  of  nat- 
ural waters,  but  even  this  contains  air,  dust,  and  not  infre- 
quently estimable  amounts  of  nitrites  and  nitrates  of  ammonium. 
All  water  that  has  been  in  contact  with  the  soil  contains  some 
of  the  ingredients  of  the  latter  in  solution.  On  evaporating 


WATER 


37 


off  the  water,  these  dissolved  ingredients,  which  are  in  the 
main  salts  of  various  kinds,  are  left  behind  as  a  residue.  The 
amount  of  material  taken  up  from  the  soil  by  water  varies  very 
greatly  with  the  nature  of  the  soil.  Thus  from  soil  formed 
mainly  from  the  disintegration  of  granite  rocks,  relatively  small 
amounts  of  material  are  dissolved,  whereas,  from  limestone  soils 
large  quantities  enter  into  solution.  By  distilling  natural 
waters,  they  may  be  freed  from  the  dissolved,  non- volatile  in- 
gredients. In  this  way  pure  water  may  be  obtained.  The 
process  consists  of  boiling  the  water  in  a  retort  and  condensing 
the  steam  formed  (Fig.  20).  In  this  process  the  condenser  is, 


FIG.  20. 


of  course,  always  dissolved  to  a  slight  extent.  The  material 
of  which  it  is  constructed  is  somewhat  soluble  in  water.  Thus 
glass  condensers  are  always  somewhat  attacked  by  water, 
though  not  sufficiently  so  to  make  the  distilled  water  unfit  for 
ordinary  purposes.  When  a  very  pure  water  is  desired,  a  block 
tin,  or,  still  better,  a  platinum  condenser,  is  used.  On  boiling 
water,  the  dissolved  gases  it  contains  are  almost  completely  ex- 
pelled. Distilled  water  tastes  flat ;  whereas  natural  waters, 
which  contain  air,  have  a  refreshing  taste. 

Natural  Waters.  —  The  solid  ingredients  in  natural  waters 
vary  greatly  in  character  and  amount.  In  oceanic  waters  there 
is  about  3.5  per  cent  of  solid  material,  of  which  2.7  per  cent 
consists  of  common  salt,  and  the  remainder  mainly  of  chlorides 
and  sulphates  of  magnesium,  calcium,  and  potassium,  together 
with  smaller  amounts  of  the  bromides  and  carbonates  of  these 


38  OUTLINES  OF  CHEMISTRY 

metals.  Some  thirty  elements  occur  in  oceanic  waters,"  most 
of  them  in  very  minute  amounts.  The  water  of  the  Dead  Sea 
contains  22.8  per  cent  of  saline  matter  and  that  of  Great  Salt 
Lake  in  Utah  23.04  per  cent.  Fresh  water  as  we  find  it  in 
rivers  and  many  lakes  usually  contains  from  0.005  to  0.15 
per  cent  of  solid  material,  and  deep  well  water  averages  from 
0.01  to  0.4  per  cent.  The  amount  of  salts  contained  in  the 
waters  of  springs  and  wells  varies  greatly  with  the  character 
.  of  the  strata  of  the  earth's  crust  with  which  the  water  has  been 
in  contact.  t Sandstone  and  granitic  material  is  less  attacked 
by  water  than  soils  rich  in  the  carbonates  of  lime  and  magnesia; 
consequently,  springs  and  wells  in  limestone  regions  contain 
much  more  solid  material  in  solution  than  those  where  sand- 
stone and  granitic  rocks  abound.  Rain  water  is  really  distilled 
water  ;  though  as  it  falls  through  the  atmosphere  it  gathers 
dust  and  dissolves  the  atmospheric  gases.  If  water  is  gathered 
during  a  shower,  that  which  falls  after  a  time  is  much  purer 
than  that  which  first  falls  to  earth.  This  is  due  to  the  fact 
that  the  air  is  fairly  well  washed  during  the  earlier  part  of  a 
copious  rainfall.  Waters  containing  a  large  amount  of  calcium 
salts  are  called  hard  waters.  They  do  not  form  a  lather  with 
soap,  and  do  not  soften  vegetables  properly  when  these  are 
boiled  in  such  water.  Furthermore,  these  waters  produce  a 
hard  sediment  consisting  mainly  of  carbonate  of  lime  which 
clogs  up  cooking  utensils,  'boilers,  and  pipes.  The  purification 
of  hard  waters  will  be  considered  in  connection  with  the  salts 
of  calcium. 

Potable  Water.  —  For  ordinary  drinking  purposes,  water 
should  be  colorless,  odorless,  tasteless,  and  free  from  materials 
that  may  prove  to  be  deleterious  to  health.  The  mineral  in- 
gredients commonly  found  in  waters  from  springs,  wells,  brooks, 
rivers,  and  lakes  are  not  injurious  to  the  system.  It  is  when 
these  sources  are  contaminated  by  sewage,  which  very  fre- 
quently gets  into  them,  that  the  waters  become  dangerous  to 
health  ;  for  the  organic  animal  and  vegetable  material  in  de- 
composing develops  products  which  may  be  injurious,  and  often 
affords  a  place  for  the  growth  of  bacteria  that  cause  disease. 
For  this  reason,  any  water  that  clearly  shows  that  it  has.  been 
contaminated  by  sewage  is  pronounced  dangerous  to  health. 

It  is  clear  that  a  bacteriological  examination  ought  to  accom- 


WATER  39 

pany  a  chemical  examination  of  a  drinking  water,  for  injurious 
organisms  may  be  present  in  water  even  though  the  sewage 
contamination  be  so  slight  that  a  chemist  would  pronounce  the 
water  fit  to  drink.  As  common  salt  and  organic  matter  and 
its  decomposition  products,  especially  nitrites  and  nitrates  of 
ammonium,  characterize  sewage,  the  determination  of  the 
amounts  of  these  ingredients  forms  the  chief  task  of  the  chem- 
ist in  analyzing  a  potable  water.  The  air  dissolved  in  natural 
waters  renders  it  refreshing.  As  boiling  kills  the  disease  germs 
in  water,  it  is  frequently  resorted  to,  especially  in  large  cities, 
in  cases  of  epidemics  caused  by  contaminated  water.  The 
process  of  boiling  expels  the  gases  dissolved  in  the  water  and 
renders  it  insipid  to  the  taste.  Thus  wholesome  drinking 
water  is  not  at  all  chemically  pure  water.  The  latter  is  not 
even  common  in  chemical  laboratories,  for  ordinary  distilled 
water,  though  free  from  non- volatile  ingredients,  still  contains 
carbon  dioxide,  air,  and  not  infrequently  ammonium  salts  in 
solution.  These  impurities  are  not  of  consequence,  however, 
for  ordinary  purposes. 

Very  frequently,  river  and  lake  water  must  be  used  for 
drinking  purposes,  even  though  it  is  somewhat  contaminated 
by  sewage.  These  waters  must  then  be  subjected  to  purifica- 
tion, which  commonly  consists  of  filtration  through  beds  of 
sand  and  exposure  to  the  air,  the  oxygen  of  which  being  ab- 
sorbed by  the  water,  oxidizes  the  organic  material  to  simpler 
products  that  are  comparatively  harmless  to  the  human  system. 
The  filters,  of  course,  must  be  renewed  from  time  to  time,  for 
the  organic  material  collects  in  them  and  thus  they  may  after 
a  while  themselves  become  a  source  of  contamination.  On  a 
small  scale,  the  Pasteur-Chamberland  water  filter  is  entirely 
efficient  in  freeing  water  from  suspended  matter  and  bacteria. 
This  filter  consists  of  unglazed  porcelain,  generally  in  form  of  a 
tube  closed  on  one  end,  which  is  attached  to  the  ordinary  water 
cock.  The  water  thus  filters  through  the  pores  of  this  un- 
glazed porcelain  under  the  pressure  of  the  waterworks  system. 

Mineral  Water.  —  Waters  containing  special  mineral  ingre- 
dients or  dissolved  gases  are  frequently  used  for  medicinal  pur- 
poses. Among  the  mineral  waters  are  distinguished  :  (1)  bit- 
ter waters  that  are  rich  in  magnesium  salts  ;  (2)  chalybeate 
waters  that  contain  iron  salts  ;  (3)  sulphur  waters  which  con- 


40  OUTLINES  OF   CHEMISTRY 

tain  hydrogen  sulphide  ;  (4)  carbonated  waters  which  are 
charged  with  carbon  dioxide  so  that  they  effervesce  ;  (5)  lithia 
waters  containing  lithium  salts.  Thermal  waters  are  those 
which  have  a  higher  temperature  than  the  surrounding  atmos- 
phere. They  frequently  also  contain  special  mineral  ingredi- 
ents which  are  considered  valuable  for  therapeutic  purposes. 

Composition.  —  Chemically  pure  water  is  a  compound  of 
oxygen  and  hydrogen,  two  volumes  of  the  latter  uniting  with 
one  volume  of  the  former  to  form  water.  By  weight  water 
consists  of  88.864  per  cent  oxygen  and  11.136  per  cent  hydro- 
gen. Knowing  that  by  the  electrolysis  of  water  two  volumes 
of  hydrogen  and  one  volume  of  oxygen  are  produced,  and  hav- 
ing found  the  weight  of  a  liter  of  hydrogen  and  that  of  a  liter 
of  oxygen,  the  composition  of  water  by  weight  can  readily  be 
computed. 

When  hydrogen  is  passed  over  copper  oxide  heated  to  a  dull 
redness,  the  oxide  is  reduced  to  metallic  copper  and  water  is 
formed.  Consequently,  by  heating  a  known  amount  of  dry 
copper  oxide  in  a  tube,  in  a  current  of  dry  hydrogen,  and  col- 
lecting and  weighing  the  water  formed,  and  also  weighing  the 
metallic  copper  obtained,  the  percentage  composition  of  water 
may  be  computed.  Obviously,  the  loss  of  weight  of  the  copper 
oxide  represents  the  oxygen  that  has  entered  into  combination 
in  the  water  formed  ;  and  the  difference  between  the  weight 
of  the  latter  and  the  oxygen  given  off  by  the  copper  oxide 
is  the  weight  of  the  hydrogen  in  the  water  produced.  This 
method  of  determining  the  composition  of  water  was  used  by 
Dulong  and  Berzelius  in  1819.  It  was  also  employed  by  Dumas 
in  1842  with  greater  refinements. 

The  researches  on  the  composition  of  water  have  yielded  the 
result  that  for  each  gram  of  hydrogen,  water  contains  7.94 
grams  of  oxygen;  that  is  to  say,  the  ratio  of  hydrogen  to 
oxygen  in  water  is  nearly  1  to  8. 

Gay-Lussac's  Law  of  Combination  of  Gases  by  Volume.  — 
When  two  volumes  of  hydrogen  and  one  volume  of  oxygen  unite 
chemically,  and  the  water  formed  is  not  allowed  to  condense  to 
the  liquid  state,  it  is  found  that  the  steam  obtained  occupies 
two  volumes,  measured,  of  course,  at  the  same  temperature  and 
pressure  as  the  oxygen  and  hydrogen.  To  demonstrate  this, 
the  apparatus  of  Hoffman  (Fig.  21)  is  convenient.  The  inner 


WATER 


41 


long  eudiometer  tube  A  is  filled  with  mercury  and  then  in- 
verted in  the  mercury  bath  B.  Thus,  a  Torricelli  vacuum  is 
formed  in  A,  whose  upper  end  is  provided  with  a  pair  of  plati- 


FIG.  21. 


42  OUTLINES   OF  CHEMISTRY 

num  terminals,  across  which  an  electric  spark  may  be  passed 
by  connecting  with  the  induction  coil  0.  The  eudiometer 
tube  A  is  placed  inside  of  the  larger  tube  D,  which  is  filled 
with  steam  from  the  boiler  E.  By  this  means  the  eudiometer 
tube  is  heated  to  the  boiling  point  of  water.  If  now  a  mixture 
of  two  volumes  of  hydrogen  and  one  volume  of  oxygen  is  intro- 
duced into  the  eudiometer  tube  A,  the  volume  carefully  noted, 
and  then  the  mixture  is  exploded  by  passing  the  electric 
spark,  the  resulting  water  vapor  will  be  found  to  have  two 
thirds  of  the  volume  of  the  mixture  of  the  oxy hydrogen  gas 
introduced,  when  the  level  of  the  mercury  in  the  eudiometer 
has  been  restored  to  its  original  place.  Therefore,  two  volumes 
of  hydrogen  and  one  volume  of  oxygen  unite  to  form  two  volumes 
of  water  vapor.  This  very  simple  relation  is  typical  of  the 
volume  relations  in  general  which  have  been  found  to  obtain 
when  gases  combine  chemically.  Expressed  in  general  terms 
we  may  say  :  When  gases  combine  chemically  with  one  another, 
the  volumes  that  unite  bear  a  simple  relation  to  one  another  ;  and 
if  the  product  formed  be  gaseous,  its  volume  also  bears  a  simple 
relation  to  the  volumes  of  the  original  gases  that  have  entered  into 
combination.  This  law  was  discovered  by  Joseph  Louis  Gay- 
Lussac,  professor  of  chemistry  at  the  Sorbonne,  Paris.  We 
shall  meet  with  further  specific  illustrations  of  this  law,  which 
is  known  as  the  law  of  Gay-Lussac  of  combination  of  gases  by 
volume.  It  is  of  great  importance  in  chemistry,  as  will  appear 
in  the  succeeding  chapters. 

Properties  of  Water.  —  In  thin  layers,  pure  water  is  colorless, 
but  in  deep  layers  it  has  a  greenish  blue  color.  This  explains 
the  beautiful  hue  of  the  waters  of  the  sea  and  many  lakes. 
River  waters  are  commonly  brownish  in  color,  due  to  the  humus 
material  which  they  contain  from  the  soils  through  which  they 
have  coursed.  The  freezing  point  of  water  is  taken  as  the  zero 
of  the  centigrade  scale,  and  the  boiling  point  under  a  pressure 
of  76  cm.  of  mercury  is  taken  as  the  100°  point  of  that  scale. 
At  and  below  360°  C  water  may  be  condensed  to  a  liquid ; 
above  this  point,  which  is  the  critical  temperature,  water  is  a 
gas  which  cannot  be  condensed  to  a  liquid  even  though  very 
high  pressures  be  applied. 

Like  liquids  in  general,  water  is  but  slightly  compressible. 
Thus,  by  placing  a  liter  of  water  at  20°  under  a  pressure  of  two 


WATER  43 

atmospheres  its  volume  is  diminished  only  by  0.046  of  a  cubic 
centimeter.  The  volume  of  a  given  weight  of  water  varies 
with  the  temperature.  Water  expands  in  volume  when  heated 
above  4°,  and  also  when  cooled  below  that  temperature  to  its 
freezing  point.  Water,  therefore,  has  its  maximum  density  at 
4°.  Most  substances  show  a  continuous  diminution  in  volume 
when  cooled.  The  fact  that  water  expands  when  cooled  below 
4°  is  therefore  a  very  exceptional  behavior.  At  4°  a  cubic 
centimeter  of  water  weighs  one  gram.  Water  at  its  maximum 
density  is  commonly  taken  as  the  standard  liquid  with  which 
the  weights  of  equal  volumes  of  other  liquids  and  solids  are 
compared.  In  other  words,  water  at  4°  is  the  standard  of 
comparison  of  the  specific  gravities  of  liquids  and  solids. 
At  100°  the  volume  of  water  is  about  4.3  per  cent  greater 
than  at  0°. 

The  amount  of  heat  required  to  raise  the  temperature  of  one 
gram  of  water  one  degree  is  called  a  calorie  (cal.) ;  it  is  the  unit 
used  in  the  measurement  of  heat.  It  requires  80  cal.  to  trans- 
form one  gram  of  ice  at  0°  to  water  of  the  same  temperature ; 
i.e.  the  latent  heat  of  fusion  of  ice  is  80  cal.  To  convert  one 
gram  of  water  at  100°  into  vapor  of  the  same  temperature  re- 
quires 537  cal.,  which  is  the  so-called  latent  heat  of  evapora- 
tion of  water.  The  specific  heat,  the  latent  heat  of  fusion, 
and  the  latent  heat  of  evaporation  of  water  are  very  high 
indeed,  as  compared  with  similar  constants  of  most  other 
substances. 

When  water  freezes  it  expands,  and  the  ice  at  0°  occupies 
1.0908  times  the  volume  of  the  water  at  the  same  temperature. 
This  behavior  of  water  is  again  unusual,  for  most  substances 
contract  during  the  act  of  congealing,  thus  forming  a  solid  that 
is  heavier  than  the  liquid.  The  fact  that  water  increases  in 
volume  as  it  solidifies  is  an  important  factor  in  the  disintegra- 
tion of  rocks,  for  the  force  exerted  by  water  in  freezing  is 
enormous.  The  bursting  of  frozen  water  pipes  and  other  con- 
tainers in  winter  is  also  due  to  the  expansion  of  water  in  freez- 
ing. But  the  fact  that  ice  is  lighter  than  water  is  of  further 
importance  in  nature ;  for  were  it  not  for  this,  many  of  our  lakes 
and  rivers  would  freeze  to  the  bottom  in  winter,  and  thus  fishes 
and  other  organisms  in  these  waters  would  be  destroyed.  The 
huge  masses  of  ice  that  would  accumulate  in  winter  also  would 


44  OUTLINES  OF   CHEMISTRY 

materially  reduce  the  temperature  for  the  remainder  of  the 
year. 

Supercooled  Water.  —  Ice  melts  at  0°,  but  when  water  is  cooled 
to  0°,  it  does  not  necessarily  freeze.  In  fact,  water  may  be 
cooled  several  degrees  below  zero  and  still  be  liquid.  Water 
in  this  condition  is  said  to  be  supercooled,  or  in  a  metastable 
condition.  If  water  thus  supercooled  is  brought  in  contact 
with  a  piece  of  ice,  the  whole  mass  freezes  to  a  solid.  If  super- 
cooled water  is  cooled  still  further,  a  point  is  finally  reached  at 
which  it  will  congeal  without  being  touched  with  ice.  Super- 
cooled water  may  be  kept  in  the  liquid  condition  for  a  long 
time.  Sometimes  shaking,  jarring,  or  brisk  stirring  induces 
freezing  of  supercooled  water,  but  this  is  not  necessarily  the 
case.  The  lower  the  temperature  of  the  metastable  water,  the 
more  likely  is  jarring  to  induce  ice  formation.  However,  touch- 
ing supercooled  water  with  ice,  always  causes  freezing. 

The  freezing  point  and  the  melting  point  of  water  are  the 
same ;  namely,  0°.  This  is  the  temperature  at  which  ice  and 
water  are  in  equilibrium  with  each  other  at  ordinary  pressure. 
Raise  the  temperature  above  0°  and  all  ice  disappears  ;  cool 
below  0°  in  the  presence  of  ice  and  the  whole  mass  freezes;  i.e. 
liquid  water  disappears.  Similarly,  the  freezing  or  melting  point 
of  any  solid  is  an  equilibrium  temperature  at  which  the  solid  and 
liquid  can  exist  side  by  side  in  contact  with  each  other  without 
change. 

Change  of  Freezing  Point  with  Pressure.  —  The  freezing 
point  of  ice  is  altered  by  change  of  pressure.  Since  water  ex- 
pands on  congealing,  an  increase  of  pressure  on  its  surface 
would  make  it  more  difficult  for  ice  to  form.  In  other  words, 
we  should  then  have  to  cool  water  under  pressure  to  a  lower 
temperature  in  order  to  freeze  it ;  or  what  comes  to  the  same 
thing,  ice  under  pressure  melts  at  a  lower  temperature  than  at 
ordinary  pressure.  Substances  which  do  not  expand,  but  con- 
tract as  they  congeal,  act  just  the  opposite  from  water  in  this 
respect  when  put  under  pressure  ;  i.e.  increase  of  pressure 
causes  them  to  freeze  at  a  higher  temperature,  the  increase  of 
pressure  aiding  contraction  which  accompanies  the  solidification 
in  these  cases. 

These  instances  of  the  alteration  of  the  freezing  point  of 
substances  with  increase  of  pressure  are  illustrations  of  a  far- 


WATER  45 

reaching  principle  which  may  be  stated  as  follows  :  When 
chemical  or  physical  equilibrium  exists,  and  one  of  the  factors 
upon  which  it  depends  is  altered,  a  change  is  produced  which 
opposes  the  first  alteration. 

This  is  known  as  the  principle  of  Le  Chatelier,  who  first  enun- 
ciated it.  Thus  increase  of  pressure  upon  any  solid  or  liquid 
tends  to  diminish  its  volume.  When  ice  and  water  exist  in 
equilibrium  at  0°  and  the  pressure  is  increased,  the  ice  melts, 
which  process  is  accompanied  with  a  diminution  in  volume, 
which  has  a  tendency  to  lessen  the  pressure.  In  the  case  of  in- 
crease of  pressure  upon  solid  and  liquid  tin  in  equilibrium 
with  each  other  at  the  melting  point  of  tin,  the  liquid  tin  will 
congeal,  for  thus  contraction  is  brought  about  and  consequently 
the  pressure  exerted  upon  the  tin  is  lessened.  The  principle 
of  Le  Chatelier  is  of  far-reaching  importance,  and  we  shall  have 
further  examples  of  it  later. 

Crystalline  Nature  of  Ice.  —  When  water  solidifies,  it  tends 
to  take  on  regular  forms.  This  is  evident  from  the  frost  on 
the  windows,  from  the  shapes  of  snowflakes,  and  the  radial 
structure  of  ice.  The  needles  that  form  as  ice  congeals  tend  to 
arrange  themselves  so  as  to  form  angles  of  60°.  These  forms 
are  most  perfect  perhaps  in  the  case  of  snowflakes,  which 
as  they  fall  on  a  still  day  are  frequently  quite  large  and 
perfect. 

Water  crystallizes  in  the  hexagonal  system,  which  is  one  of 
the  six  systems  into  which  all  known  crystals  may  be  divided 
(see  Crystal  Systems).  Not  only  do  crystals  exhibit  outward 
regularity  of  form,  but  they  also  show  different  degrees  of 
hardness,  tenacity,  refrangibility,  light  absorbing  power,  etc., 
in  different  directions.  We  therefore  distinguish  crystalline 
substances,  which  show  these  characteristics,  from  amorphous 
substances,  which  do  not  have  regularity  of  form  and  which 
exhibit  the  same  properties  irrespective  of  the  direction  through 
them.  Ice  is  a  typical  crystalline  substance,  while  glass  is 
a  typical  amorphous  substance.  Amorphous  means  without 
form.  Crystalline  substances  commonly  have  a  definite  melt- 
ing point  and  definite  solubility,  while  amorphous  substances 
do  not. 

Thus  glass  has  no  definite  temperature  at  which  it  melts. 
It  softens  when  heated  and  gradually  passes  through  all  stages 


46 


OUTLINES   OF  CHEMISTRY 


of  gradations  to  the  liquid  state  on  further  heating.  Not  so 
with  water,  for  it  has  a  sharp  melting  point  at  0°.  Many 
definite  chemical  compounds  tend  to  form  crystals;  and  since 
the  same  compound  tends  to  take  on  the  same  shape  under 
given  conditions,  the  study  of  crystallography  is  of  value  to 
the  chemist  in  aiding  him  in  purifying  and  identifying  sub- 
stances. However,  many  definite  chemical  compounds  have 
never  been  obtained  in  the  crystalline  condition. 

Compounds  with  Water.  —  Many  salts,  like  copper  sulphate, 
Glauber's  salt,  and  Epsom  salt,  crystallize  with  water.  The 
water  in  these  salts  is  spoken  of  as  water  of  crystallization. 
On  exposure  to  the  air,  some  of  these  salts  lose  a  portion  of 
this  water  of  crystallization  and  become  opaque  or  crumble  to 
a  powder.  They  are  said  to  effloresce.  Other  salts,  like  calcium 
chloride,  have  such  a  strong  attraction  for  water,  that  on  ex- 
posure to  the  air  they  take  on  water  from  the  air  and  finally 
become  completely  dissolved.  They  are  said  to  deliquesce. 
Substances  that  have  attraction  for  water  are  also  called  hygro- 
scopic. Concentrated  sulphuric  acid,  phosphorus  pentoxide, 

calcium  chloride,  lime,  and 
caustic  potash  are  strongly 
hygroscopic.  Gases  passed 
over  these  are  deprived  of 
their  moisture,  and  many 
solids  left  with  them  for  a 
time  in  a  confined  space  are 
dried  or  desiccated.  A  typical 
form  of  desiccator  is  shown  in 
Fig,  22.  The  strongly  hy- 
groscopic substance  is  placed 
in  the  bottom  of  the  vessel, 
and  the  substance  to  be  dried 

is  placed  on  the  support  in  the  upper  part  of  the  apparatus. 
Through  the  cock  the  air  can  be  exhausted  from  the  apparatus 
and  thus  the  drying  process  be  aided  still  further.  Such 
desiccators  are  frequently  used  in  chemical  work,  for  many 
substances  like  glass,  porcelain,  and  even  metallic  utensils 
attract  moisture  and  form  a  film  of  it  on  their  surface.  This 
film  varies  in  thickness  and  weight  with  the  degree  of  humidity 
of  the  atmosphere.  In  accurate  quantitative  experiments  it  is 


FIG.  22. 


WATER  47 

necessary  to  eliminate  this  film  of  moisture,  and  for  this  purpose 
desiccators  are  commonly  used.  If  permissible,  the  objects  are 
heated  to  drive  off  moisture,  and  then  cooled  in  the  desiccator. 
If  heating  is  not  permissible  in  a  given  case,  the  substance  is 
introduced  into  the  desiccator  and  kept  there  for  a  longer 
time,  frequently  in  a  vacuum.  It  is  evident  that  in  a  desic- 
cator, the  drying  material  used  must  have  a  greater  affinity  for 
water  than  the  substance  to  be  dried. 

When  water  simply  adds  itself  to  another  compound,  the 
product  is  commonly  termed  a  hydrate.  Such  hydrates  are 
quite  common;  thus  ferric  chloride  forms  several  hydrates 
with  water,  which  follow  the  laws  of  definite  and  multiple 
proportions. 

When  oxides  unite  with  water,  or  when  a  metal  like  sodium 
acts  on  water,  crowding  out  a  portion  of  its  hydrogen,  the 
product  formed  is  commonly  termed  an  hydroxide.  In  these 
cases  it  is  always  possible  to  regard  the  resulting  substance  as 
water  in  which  a  portion  of  the  hydrogen  has  been  replaced  by 
another  element.  So  when  lime  and  water  act  on  each  other 
they  unite  and  form  slaked  lime,  which  is  hydroxide  of  calcium. 
Caustic  potash,  which  may  be  formed  by  the  action  of  potas- 
sium on  water,  with  concomitant  evolution  of  hydrogen,  is 
potassium  hydroxide. 

Water  as  a  Solvent.  —  Many  substances  are  soluble  in  water. 
Indeed,  of  so  many  is  this  the  case  that  water  has  at  times  been 
termed  a  universal  solvent.  There  are,  however,  very  many 
compounds  that  are  not  soluble  in  water.  In  general,  the 
ordinary  acids,  alkalies,  and  salts  used  in  the  chemical  lab- 
oratory are  soluble  in  water  to  a  greater  or  lesser  degree. 
The  rocks  of  the  earth's  crust  are  all  soluble  to  some  extent, 
though  to  a  very  slight  degree  in  some  cases ;  yet  this  slight 
solubility  of  rocks  is  of  the  highest  importance  to  plants  whose 
rootlets  are  thus  able  to  take  up  mineral  matter  needed  for 
their  economy  and  growth.  In  geological  transformations, 
such  as  the  weathering  of  rocks,  the  formation  of  soils,  and  the 
deposition  of  ores,  this  slight  solubility  is  nevertheless  the 
determining  factor,  without  which  these  processes  could  not 
proceed. 

Fats,  waxes,  oils,  and  kindred  substances  are  generally  not 
soluble  in  water.  Yet  many  of  these  have  some  degree  of 


48  OUTLINES  OF  CHEMISTRY 

attraction  for  water,  which  is  shown  by  the  fact  that  they  are 
often  slightly  hygroscopic.  And  again,  in  the  bodies  of  plants, 
and  particularly  in  those  of  animals,  fatty  material  is  very 
closely  associated  with  tissues  which  are  rich  in  water.  So 
that  although  fats  are  generally  not  soluble  in  water  to  speak 
of,  yet  in  many  cases  there  is  evidence  that  some  degree  of 
attraction  between  them  and  water  does  exist.  Solutions  will 
receive  further  consideration  later. 


CHAPTER  V 

HYDROCHLORIC  ACID  AND   CHLORINE 

Preparation  and  Properties  of  Hydrochloric  Acid.  —  When  con- 
centrated sulphuric  acid  is  poured  upon  common  salt,  an  effer- 
vescence ensues,  a  gas  being  evolved  which  is  colorless,  has  a 
pungent  odor,  and  is  neither  combustible  nor  a  supporter  of 
combustion.  This  gas  has  a  very  sour  taste,  and  produces 
suffocation  when  inhaled  in  quantity.  It  reddens  moist  blue 
litmus  paper,  and  is  very  soluble  in  water.  At  0°  one  volume 
of  water  will  absorb  503  volumes  of  the  gas,  while  at  room  tem- 
perature about  450  volumes  are  thus  absorbed.  This  gas,  which 
was  at  first  called  "  spirit  of  salt,"  was  discovered  by  Johann 
Rudolf  Glauber  in  1658.  It  is  hydrochloric  acid.  Priestley 
called  it  "  marine  acid  air  "  ;  he  collected  the  gas  over  mercury. 

Hydrochloric  acid  is  sometimes  emitted  during  volcanic  erup- 
tions. It  also  occurs  in  the  gastric  juice  of  man  and  other 
animals.  In  normal  condition  the  human  gastric  juice  contains 
about  0.33  per  cent  of  hydrochloric  acid,  which  is  essential  in 
the  process  of  digestion. 

Hydrochloric  acid  comes  in  the  market  as  a  solution  of  the 
gas  in  water.  It  also  goes  by  the  name  of  muriatic  acid.  On 
dissolving  pure  hydrochloric  acid  gas  in  distilled  water,  a  color- 
less solution  is  obtained.  However,  much  of  the  commercial 
muriatic  acid  is  colored  yellowish  by  the  impurities,  especially 
salts  of  iron,  that  it  contains. 

The  attraction  between  hydrochloric  acid  gas  and  water  is  so 
great  that  the  gas  fumes  strongly  in  the  air.  This  is  due  to 
the  fact  that  it  condenses  moisture  from  the  air  in  drops,  which 
consist  of  an  aqueous  hydrochloric  acid  solution.  When  the 
gas  is  conducted  into  water,  heat  is  evolved.  The  thermal 
change  accompanying  the  solution  of  any  substance  is  termed 
the  heat  of  solution  (see  Thermochemistry).  Aqueous  solutions 
of  hydrochloric  acid  are  heavier  than  water.  Thus,  a  solution 
of  specific  gravity  1.024  at  15°  contains  5  per  cent  hydrochloric 
E  49 


50 


OUTLINES  OF  CHEMISTRY 


acid,  while  solutions  having  the  specific  gravities  1.049,  1.100, 
1.152,  and  1.200  contain  10,  20,  30,  and  40  per  cent,  respec- 
tively. A  solution  which  is  saturated  with  hydrochloric  acid 
at  15°  contains  42.9  per  cent  and  has  a  specific  gravity  of  1.212. 
The  usual  "  pure,"  commercial,  concentrated  hydrochloric  acid 
is  about  38  per  cent  strong  and  has  a  specific  gravity  of  1.19. 
It  fumes  strongly  when  exposed  to  the  air. 

On  boiling  a  saturated  solution  of  hydrochloric  acid,  the  gas 
is  in  part  expelled,  and  finally  a  20.2  per  cent  solution  with  a 
boiling  point  of  110°  is  obtained.  At  ordinary  pressure,  this 
solution  distills  over  without  change  of  composition.  The 
same  strength  of  solution  is  finally  obtained  when  a  dilute  solu- 
tion is  boiled.  In  this  case  water  is  mainly  expelled  until  the 
solution  reaches  a  strength  of  20.2  per  cent,  when  it  distills 
over  without  decomposition.  The  final  acid  thus  obtained  at 
different  pressures,  however,  has  a  slightly 
different  composition. 

Pure,  dry  hydrochloric  acid  gas  may  be 
condensed  to  a  liquid  at  10°  under  a  pres- 
sure of  40  atmospheres.  Under  atmospheric 
pressure  the  liquid,  which  is  colorless,  boils 
at  —  84°  and  freezes  at  about  -  110°. 

Composition  and  Chemical  Behavior  of 
Hydrochloric  Acid.  —  When  metallic  sodium 
is  introduced  into  pure  hydrochloric  acid 
gas,  the  metal  burns  in  the  gas,  forming 
common  salt  and  hydrogen.  This  fact 
shows  that  hydrogen  is  one  of  the  con- 
stituents of  hydrochloric  acid. 

The  right  limb  of  the  apparatus  (Fig. 
23)  is  filled  with  pure,  dry  hydrochloric 
acid  gas.  The  press  P,  which  fits  securely 
on  the  top  of  the  glass  tube,  contains 
metallic  sodium.  When  the  latter  metal 
is  pressed  out  into  the  tube  A,  by  turning 
the  screw  of  the  press,  the  sodium  and 
hydrochloric  acid  react  and  form  common 
salt  and  hydrogen  with  concomitant  evolu- 
tion of  light  and  heat.  If  the  level  in  the 
FIG.  23.  limbs  A  and  B  is  kept  constant  by  pouring 


HYDROCHLORIC   ACID  AND  CHLORINE 


51 


mercury  into  B  as  required,  it  will  be  seen  that  wheu  further 
addition  of  sodium  no  longer  produces  any  action  in  A,  the 
hydrogen  obtained  occupies  just  one  half  of  the  vocume  of  the 
original  hydrochloric  acid  gas. 

Hydrochloric  acid  is  a  very  powerful  acid  and  acts  strongly 
on  many  metals,  hydrogen  being  liberated  and  chlorides  of  the 
metals  being  formed  during  the  reaction. 

When  a  concentrated  aqueous  solution  of  hydrochloric  acid  is 
subjected  to  electrolysis  {Fig.  24),  equal  volumes  of  hydrogen  and 


FIG.  24. 

a  greenish  yellow  gas,  chlorine,  appear.  Carbon  electrodes  are 
used  in  this  electrolysis,  for  platinum  would  be  attacked  by  the 
chlorine.  This  apparatus,  designed  by  Lothar  Meyer,  differs 
from  that  in  Fig.  2,  because  chlorine  when  collected  over  an 
aqueous  hydrochloric  acid  solution  under  pressure  is  quite 
appreciably  absorbed,  so  that  the  volume  of  the  chlorine  gas 
would  be  diminished. 

When  equal  volumes  of  dry  chlorine  and  dry  hydrogen  con- 
tained in  the  two  parts  of  the  strong  tube  (Fig.  25)  are  allowed 
to  mix  by  opening  the  stopcock,  and  the  mixture  is  then  ex- 
posed to  diffused  daylight,  hydrochloric  acid  is  formed,  and 
neither  hydrogen  nor  chlorine  is  left  uncombined.  Moreover, 


52 


OUTLINES  OF  CHEMISTRY 


the  volume  of  the  hydrochloric  acid  gas  formed  is  exactly 
equal  to  that  of  the  hydrogen  plus  chlorine.  That  is,  equal 
volumes  of  hydrogen  and  chlorine  unite  to  form  hydrochloric  acid 

without  change  of  volume,  which 
is    demonstrated    by   the    fact 
that  when  the  stopper  at  the 
lower  end   of   the  tube   (Fig. 
25)  is  removed  under  mercury 
after  the  hydrochloric  acid  has 
formed,  neither  gas  escapes  nor 
mercury  enters   the  tube.     In 
direct  sunlight  or  on  exposure 
to  a   strong   magnesium   flash 
light    the    union    takes    place 
with  explosive  violence.     Thus 
it  is  that  one  volume  of  hydrogen 
unites  with  one  volume  of  chlorine 
to  form  two  volumes  of  hydro- 
chloric acid  gas.     This   is  an- 
other example  illustrating  the 
law  of  Gay-Lussac  of  combina- 
tion of  gases  by  volume. 
It  has  been  found  that  one  volume  of  chlorine  is  35.45  times 
as  heavy  as  an  equal  volume  of  hydrogen.     From  this  and  the 
fact  that  equal  volumes  of  hydrogen  and  chlorine  unite  to  form 
hydrochloric  acid,  it  is  evident  that  ly  weight,  1  part  of  hydro- 
gen unites  with  35.45  parts  of  chlorine  to  form  hydrochloric  acid. 
According  to  H.  Sainte-Claire  Deville,  hydrochloric  acid  gas 
is  partially  decomposed  into  hydrogen  and  chlorine  when  heated 
to  temperatures  of  1300°  or  above. 

Enormous  quantities  of  hydrochloric  acid  are  manufactured 
as  a  by-product  of  the  Le  Blanc  soda  process  (which  see). 

Occurrence,  History,  and  Preparation  of  Chlorine.  —  Chlorine 
occurs  in  nature  only  in  combination  with  other  elements.  The 
chlorine-bearing  compounds  are  chiefly  common  salt,  the  chloride 
of  sodium,  and  the  chlorides  of  potassium,  magnesium,  and  cal- 
cium. Chlorine  is  also  found  in  the  native  chlorides  of  lead, 
copper,  and  silver.  In  combination  with  hydrogen,  it  occurs 
in  the  gastric  juice,  and  as  sodium  chloride  and  potassium  chlo- 
ride it  forms  an  essential  constituent  of  the  bodies  of  animals. 


FIG.  25. 


HYDROCHLORIC   ACID   AND   CHLORINE  53 

It  is  also  an  important  constituent  of  plants,  in  which  it  is 
probably  mainly  combined  with  potassium. 

Chlorine  was  first  prepared  in  the  free  state  by  Scheele  in 
1774,  who  treated  manganese  dioxide  with  hot  hydrochloric 
acid.  He  called  the  gas  "  dephlogisticated  hydrochloric  acid," 
for  at  that  time  hydrogen  was  regarded  as  the  phlogiston  of 
Stahl.  However,  in  1785  Berthollet,  who  belonged  to  the  anti- 
phlogistic school,  called  chlorine  "oxidized  hydrochloric  acid." 
He  was  of  the  opinion  that  chlorine  contained  oxygen,  and  this 
view  prevailed  till  1807 ;  when,  on  the  basis  of  their  researches, 
Gay-Lussac  and  Thenard  showed  the  gas  to  be  a  simple  sub- 
stance, i.e.  an  element.  The  gas  was  named  chlorine  by  Sir 
Humphry  Davy.  The  name  comes  from  the  Greek,  meaning 
greenish  yellow. 

We  have  seen  that  chlorine  is  one  of  the  products  of  the  elec- 
trolysis of  hydrochloric  acid.  The  simplest  way  of  preparing 
chlorine  is  by  treating  hydrochloric  acid  with  an  oxidizing  agent, 
whose  oxygen  unites  with  the  hydrogen  of  the  hydrochloric 
acid,  thus  forming  water  and  setting  the  chlorine  free.  As 
such  an  oxidizing  agent,  manganese  dioxide  is  commonly  em- 
ployed. Chlorine  may  be  formed  by  treating  manganese  diox- 
ide with  the  aqueous  solution  of  hydrochloric  acid  and  heating 
gently ;  or  by  mixing  common  salt  with  manganese  dioxide 
and  treating  the  mixture  with  sulphuric  acid.  In  the  latter 
case,  the  sulphuric  acid  acts  on  the  sodium  chloride  forming 
hydrochloric  acid,  which  then  acts  upon  the  manganese  dioxide 
as  before.  In  these  processes  manganous  chloride  is  also  formed. 
In  place  of  manganese  dioxide,  other  oxidizing  agents,  like  po- 
tassium dichromate,  potassium  chlorate,  red  lead,  or  bleaching 
powder,  may  be  employed.  In  preparing  chlorine  by  subtract- 
ing the  hydrogen  from  the  hydrochloric  acid  by  an  oxidizing 
agent,  the  oxygen  of  the  air  may  be  employed.  By  passing  a 
mixture  of  air  and  hydrochloric  acid  at  about  400°  over  porous 
bricks  which  have  been  soaked  with  copper  sulphate  solution, 
chlorine  is  liberated.  The  method  is  called  the  Deacon  process 
and  is  used  on  a  commercial  scale.  In  this  process  cupric  chlo- 
ride is  formed,  and  this  is  decomposed  into  cuprous  chloride  and 
chlorine.  The  cuprous  chloride  is  then  again  converted  into 
cupric  chloride,  which  suffers  decomposition  as  before,  and 
so  on. 


54  OUTLINES  OF  CHEMISTRY 

Properties  of  Chlorine.  —  Chlorine  is  a  greenish  yellow  gas 
which  is  2.5  times  as  heavy  as  air.  It  has  a  very  disagreeable 
odor,  attacks  the  mucous  membranes  strongly,  giving  rise  to 
a  cough,  and  causes  death  by  suffocation.  At  0°  it  may  be 
liquefied  by  means  of  six  atmospheres  of  pressure.  The  criti- 
cal temperature  of  the  gas  is  146°,  and  the  critical  pressure  is  84 
atmospheres.  Thus,  at  ordinary  temperatures  chlorine  is  a 
condensable  vapor.  Under  atmospheric  pressure  it  becomes 
a  liquid  at  —  34°,  its  boiling  point.  Liquid  chlorine  has  a 
golden  yellow  color.  At  — 102°  it  freezes,  forming  yellow 
crystals.  Liquid  chlorine  is  now  obtainable  in  the  market  in 
lead-lined  steel  flasks  (Fig.  13).  In  this  form  it  is  shipped  for 
use  in  laboratories  and  various  industrial  plants.  . 

Chemically,  chlorine  is  a  very  active  element,  combining  at 
ordinary  temperatures  with  evolution  of  light  and  heat  with 
sodium,  phosphorus,  arsenic,  antimony, 
and  many  other  metals  when  these  are 
introduced  into  an  atmosphere  of  the  gas 
in  the  form  of  powder  or  very  thin  sheets. 
In  all  such  cases  chlorides  form  by  direct 
union  of  the  chlorine  with  the  other  ele- 
ment. An  apparatus  for  burning  arsenic 
in  chlorine  is  shown  in  Fig.  26.  The 
cork  fits  loosely.  When  the  test  tube 
containing  the  powdered  arsenic  is  raised, 
FIG.  26.  the  latter  falls  into  the  bottle  and  unites 

with  the  chlorine  with  evolution  of  light. 

Chlorine  does  not  act  directly  on  carbon  or  nitrogen ;  but 
chlorides  of  these  elements  may  be  formed  by  the  indirect 
methods  of  double  decomposition,  as  will  appear  later.  Chlo- 
rine and  hydrogen  have  a  strong  affinity  for  each  other.  A  jet 
of  hydrogen  will  burn  in  an  atmosphere  of  chlorine,  or  a  jet  of 
chlorine  in  an  atmosphere  of  hydrogen.  In  either  case  hydro- 
chloric acid  is  formed  as  the  product.  A  lighted  taper  or  gas 
jet  will  continue  to  burn  in  chlorine,  forming  hydrogen  chloride 
and  carbon,  which  forms  dense  clouds  of  soot.  Similarly,  when 
a  strip  of  filter  paper  moistened  with  turpentine  is  introduced 
into  an  atmosphere  of  chlorine,  hydrochloric  acid  is  formed, 
much  soot  escapes  in  dense  clouds,  and  the  paper  instantly 
catches  fire. 


HYDROCHLORIC  ACID  AND   CHLORINE  55 

Chlorine  is  soluble  in  water.  At  10°  1  volume  of  water 
absorbs  about  8  volumes  of  chlorine,  and  at  50°  about  1.5 
volumes.  The  solution  is  commonly  called  chlorine  water. 
When  it  is  exposed  to  sunlight, 
the  chlorine  gradually  unites 
with  the  hydrogen  of  the 
water,  forming  hydrochloric 
acid  and  oxygen.  By  filling  a 
retort  (Fig.  27)  with  chlorine 
water  and  exposing  it  to  sun- 
light, the  solution  becomes 
colorless,  hydrochloric  acid  27 

being     formed     and     oxygen 

liberated.  The  latter  gas  collects  in  the  bulb,  as  shown  in 
Fig.  27.  By  tilting  the  retort,  this  gas  may  be  brought  into 
the  neck  of  the  vessel  and  tested  with  a  glowing  splint. 

Because  chlorine  thus  unites  with  the  hydrogen  of  water  and 
sets  oxygen  free,  which  in  turn  is  capable  of  oxidizing  sub- 
stances, chlorine  is  spoken  of  as  a  powerful  oxidizing  agent. 
Upon  this  power  to  set  free  oxygen  from  water  depends  the 
bleaching  action  of  chlorine.  When  moist  flowers,  green  leaves, 
colored  cloth,  and  paper  on  which  marks  have  been  made  with 
ordinary  ink  are  introduced  into  an  atmosphere  of  chlorine, 
they  are  bleached  ;  that  is,  the  color  is  destroyed.  Moisture  is 
essential  to  have  the  bleaching  take  place;  for  the  chlorine 
unites  with  the  hydrogen  of  the  water,  forming  hydrochloric 
acid  and  setting  oxygen  free.  The  latter  then  unites  with  the 
coloring  matter  and  destroys  it..  Printer's  ink  consists  largely 
of  carbon,  which  at  ordinary  temperatures  is  not  attacked 
either  by  oxygen  or  chlorine ;  it  consequently  is  not  bleached. 
It  should  further  be  stated  that  fabrics  dyed  with  some  of 
the  aniline  dyestuffs  also  retain  their  color,  even  when  treated 
with  chlorine  water.  By  the  action  of  chlorine  on  water,  some 
hypochlorous  acid  is  always  formed. 

Other  Uses  of  Chlorine.  —  The  oxygen  liberated  when  chlorine 
acts  upon  water  as  explained  is  very  destructive  to  organic  life  ; 
for  this  reason  chlorine  is  used  as  a  disinfectant.  Fungi  and 
disease  germs  are  rapidly  destroyed  by  the  action  of  chlorine. 

Chlorine  is  also  sometimes  used  in  extracting  gold  from  its 
ores.  By  direct  union  with  the  metallic  gold,  the  chloride  of 


56  OUTLINES  OF   CHEMISTRY 

that  metal  is  formed ;    and  this  salt  being  readily  soluble  in 
water,  can  then  be  leached  out  of  the  ores. 

Some  Compounds  of  Chlorine  with  Oxygen.  —  Chlorine  does 
not  form  compounds  with  oxygen  by  direct  union  of  the  two 
gases.  However,  by  the  indirect  method  of  double  decom- 
position, compounds  of  oxygen  and  chlorine  may  be  obtained. 
These  compounds  are  gases  which  readily  decompose. 

The  compounds  of  oxygen  and  chlorine  will  be  considered 
in  Chapter  VIII.  Here  only  two  of  these  compounds  will  be 
mentioned  briefly. 

Chlorine  Monoxide.  —  When  dry  chlorine  is  passed  over  red 
oxide  of  mercury  in  the  cold,  a  pale  yellow  gas  is  formed,  which 
does  not  have  the  greenish  tint  of  the  chlorine  and  readily 
decomposes  with  explosive  violence,  even  when  moderately 
heated.  At  5°  it  may  be  condensed  to  a  liquid  of  orange- 
yellow  color,  which  readily  explodes  in  sunlight  or  on  slight 
heating,  at  times  even  on  pouring  it  from  one  dish  to  another. 
The  gas  is  soluble  in  water.  One  volume  of  water  absorbs  about 
200  volumes  of  chlorine  monoxide  gas  at  0°. 

This  substance  is  an  oxide  of  chlorine,  and  consists  of  35.45 
parts  of  chlorine  to  8  parts  of  oxygen  by  weight.  It  is  called 
chlorine  monoxide. 

Chlorine  Dioxide.  —  By  carefully  treating  powdered  potassium 
chlorate  with  concentrated  sulphuric  acid  added  in  very  small 
quantities  at  a  time,  a  heavy,  deep  yellow  gas  is  evolved  which 
has  a  very  disagreeable  odor,  attacks  mercury,  and  is  readily 
soluble  in  water.  It  is  very  unstable,  exploding  in  the  sun- 
Iight4  or  when  heated  by  means  of  the  electric  spark  or  a  hot 
iron  rod.  In  the  cold,  it  may  be  condensed  to  a  liquid  of  dark 
red  color,  which  is  of  a  highly  explosive  nature. 

This  compound  is  an  oxide  of  chlorine,  which  contains  35.45 
parts  of  chlorine  and  32  parts  of  oxygen  by  weight.  It  is 
called  chlorine  dioxide  or  chlorine  peroxide. 

Thus,  in  the  case  of  these  two  oxides  of  chlorine  we  have 
another  illustration  of  the  law  of  multiple  proportions ;  for  in 
the  monoxide  35.45  parts  of  chlorine  are  united  with  8  parts 
of  oxygen  by  weight,  whereas  in  the  peroxide  35.45  parts  of 
chlorine  are  united  with  4  times  8  parts  of  oxygen. 

The  Law  of  Reciprocal  Proportions. — We  have  learned  that 
in  water  hydrogen  and  oxygen  are  united  in  the  proportions  of 


HYDROCHLORIC   ACID   AND   CHLORINE  57 

1  part  of  hydrogen  to  8  parts  of  oxygen  by  weight.  In  hydro- 
chloric acid  1  part  of  h}rdrogen  is  united  with  35.45  parts  of 
chlorine  by  weight.  In  chlorine  monoxide  we  find  that  35.45 
parts  of  chlorine  are  united  with  8  parts  of  oxygen  by  weight ; 
and  in  chlorine  peroxide  35.45  parts  of  chlorine  are  united  with 
4  times  8  parts  of  oxygen. 

Thus  we  see  that  the  proportions  by  weight  in  which  hydrogen 
and  oxygen  combine,  and  in  which  hydrogen  and  chlorine  com- 
bine, also  determine  the  proportions  in  which  chlorine  and  oxygen 
combine  with  each  other.  This  is  an  illustration  of  a  general 
law  which  holds  in  all  chemical  combinations.  It  may  be 
stated  thus  in  general  terms :  If  three  elements,  A,  B,  and  0, 
are  able  to  unite  to  form  chemical  compounds  with  one  another, 
the  proportions  by  weight  with  which  A.  and  £  unite  to  form  the 
compound  AB,  and  the  proportions  in  which  A  and  O  unite,  also 
determine  the  proportions  in  which  B  and  O  unite  with  each 
other.  This  law  is  known  as  the  law  of  reciprocal  proportions. 
It  was  discovered  by  Jeremias  Benjamin  Richter,  and  is  one 
of  the  fundamental  laws  of  chemical  combination  by  weight. 
In  the  further  study  of  chemistry,  the  student  will  meet 
numerous  illustrations  of  this  law. 


CHAPTER  VI 

THE  LAWS  OP  COMBINING  "WEIGHTS 
AND  COMBINING  VOLUMES  AND  THE 
ATOMIC  AND  MOLECULAR  THEORIES 

Retrospect. — In  the  preceding  chapters  we  have  found  that 
certain  general  laws  regulate  the  quantities  in  which  the 
chemical  elements  combine  to  form  compounds.  The  laws 
governing  the  combination  of  the  elements  by  weight  are  as 
follows : 

(1)  The  Law  of  Definite  Proportions. — A  chemical  compound 
always   contains  the  same   elements  in  the  -  same  proportions   by 
weight.     No   matter   when,  where,  or  by  what  process  hydro- 
chloric acid,   for  example,  is  formed,  it  always  contains  only 
the   elements   hydrogen   and    chlorine   in   the   proportions    of 
1  gram  of  hydrogen  to  35.45  grams  of  chlorine.     Water  always 
consists  of  hydrogen  and  oxygen  united  in  the  proportions  of 
1  gram  of  hydrogen  to  8  grams  of  oxygen.     Common  salt  is 
made  up  of  23  parts  of  sodium  to  35.45  parts  of  chlorine  by 
weight ;  and  similarly  every  other  chemical  compound  always 
has  exactly  the  same  invariable  qualitative  and  quantitative  com- 
position.    The  law  of  definite  proportions,  it  will  be  recalled, 
was  discovered  by  Lavoisier. 

(2)  The  Law  of  Multiple  Proportions.  —  When  any  two  ele- 
ments, A  and  B,  form  more  than  one  compound  with  each  other, 
the  amounts  of  B  that  unite  with  one  and  the  same  weight  of  A  are 
simple  rational  multiples  of  one  another.     Thus  iron  and  sulphur 
form  ferrous  sulphide,  which  consists  of  28  grams  of  iron  to 
every  16  grams  of  sulphur ;  and  pyrite  or  fool's  gold,  a  native 
sulphide  of  iron,  always  contains  28  grams  of  iron  to  every  32 
(i.e.  2  times  16)  grams  of  sulphur.     Again,  in  chlorine  monox- 
ide, every  35.45  grams  of  chlorine  are  united  with  8  grams 
of  oxygen.     In  chlorine  peroxide,  every  35.45  grams  of  chlo- 
rine are  united  with  32  (i.e.  4  times  8)  grams  of  oxygen;  and 
in  chlorine  heptoxide  (which  see)  every  35.45  grams  of  chlorine 

68 


FUNDAMENTAL   LAWS   AND   THEORIES 


59 


are  combined  with  56  (i.e.  7  times  8)  grams  of  oxygen.  In 
the  oxides  of  lead  the  proportions  of  lead  and  oxygen  by  weight 
are  as  follows  :  — 

(a)  In  the  black  oxide,  Lead  :  Oxygen  :  :  1  :  0.0387 
(£)  In  the  yellow  oxide,  Lead  :  Oxygen  :  :  1  :  0.0773 
(<?)  In  the  red  oxide,  Lead  :  Oxygen  :  :  1  :  0.103 
(d)  In  the  orange  oxide,  Lead  :  Oxygen  :  :  1  :  0.1160 
(V)  In  the  brown  oxide,  Lead  :  Oxygen  :  :  1  :  0.1547 
Now,  0.0387  :  0.0773 :  0.103  :  0.1160  :  0.1547  :  :  1  :  2  :  f  :  3  :  4. 

The  compounds  of  oxygen  with  nitrogen  (which  see)  are  five 
in  number,  and  these  furnish  a  further  good  illustration  of  the 
law  of  multiple  proportions.  These  oxides  are  all  gases,  in 
which  the  proportion  of  nitrogen  to  oxygen  by  weight  is  as 
follows :  — 


(a)  In  nitrous  oxide, 

(b)  In  nitric  oxide, 

(<?)  In  nitrogen  sesquioxide, 
(cT)  In  nitrogen  dioxide, 
(e)  In  nitrogen  pentoxide, 
Now,  0.5710:1.142:1.713 


Nitrogen 
Nitrogen 
Nitrogen 
Nitrogen 
Nitrogen 


Oxygen 
Oxygen 
Oxygen 
Oxygen 
Oxygen 


0.5710 

1.142 

1.713 

2.28 

2.855 


2.28  :  2.855  :  :  1  :  2  :  3  :  4 


As  already  stated,  the  law  of  multiple  proportions  was   first 
recognized  by  John  Dalton  of  Manchester,  England. 

(3)  The  Law  of  Reciprocal  Proportions.  —  When  three  elements, 
A,  B,  and  O,  are  able  to  form  chemical  compounds  with  one 
another,  the  proportions  by  weight  in  which  A  and  B  unite  to  form 
the  compound  AB,  and  the  proportions  in  which  A  and  O  unite  to 
form  the  compound  AC,  also  determine  the  proportions  in  which 
B  and  C  unite  to  form  the  compound  BC.  Thus,  hydrogen,  oxy- 
gen, and  chlorine  are  three  elements  which  can  form  compounds 
with  one  another.  In  water  we  have,  for  every  8  grams  of  oxy- 
gen, 1  gram  of  hydrogen ;  in  hydrochloric  acid,  we  have,  for 
every  35.45  grams  of  chlorine,  1  gram  of  hydrogen;  and  in 
chlorine  monoxide  we  have,  for  every  35.45  grams  of  chlorine, 
8  grams  of  oxygen  ;  whereas  in  chlorine  peroxide  every  35.45 
grams  of  chlorine  are  united  with  4  times  8  grams  of  oxygen. 
Again,  in  sodium  chloride  every  35.45  grams  of  chlorine  are 
combined  with  23  grams  of  sodium ;  in  sodium  oxide,  every  23 
grams  of  sodium  are  united  with  8  grams  of  oxygen ;  in  chlo- 


60  OUTLINES   OF  CHEMISTRY 

rine  monoxide  every  35.45  grams  of  chlorine  are  united  with  8 
grams  of  oxygen.  Further,  take  the  elements  hydrogen,  sul- 
phur, and  oxygen  :  —  In  hydrogen  sulphide  (which  see)  every 
16  grams  of  sulphur  are  combined  with  1  gram  of  hydrogen ; 
in  water  every  8  grams  of  oxygen  are  combined  with  1  gram  of 
hydrogen ;  in  sulphur  dioxide,  which  gas  is  formed  when  sulphur 
is  burned  in  the  air  or  in  oxygen,  every  16  parts  of  sulphur  are 
united  with  16  (i.e.  2  times  8)  parts  of  oxygen.  As  mentioned 
in  the  previous  chapter,  the  law  of  reciprocal  proportions  was 
first  recognized  through  the  work  of  the  German  chemist,  Jere- 
mias  Benjamin  Richter. 

The  subject  of  the  study  of  the  weight  relations  that  obtain 
when  chemical  changes  take  place  is  called  stoichiometry,  and 
the  three  laws  just  elucidated  are  the  fundamental  laws  of 
stoichiometry. 

In  connection  with  the  study  of  hydrogen,  it  was  found  that 
23  grams  of  sodium  will  liberate  1  gram  of  hydrogen  from 
water;  32.5  grams  of  zinc  by  acting  on  dilute  hydrochloric  acid 
will  also  liberate  Igram  of  hydrogen;  9  grams  of  aluminum,  or 
12  grams  of  magnesium,  will  similarly  set  free  1  gram  of  hy- 
drogen. Therefore,  23,  32.5,  9,  and  12  are  termed  the  hydro- 
gen equivalents  of  sodium,  zinc,  aluminum,  and  magnesium, 
respectively.  Now  in  water,  which  is  an  oxide  of  hydrogen,  1 
gram  of  hydrogen  is  united  to  every  8  grams  of  oxygen;  in 
sodium  oxide,  23  grams  of  sodium  are  united  to  8  grams  of 
oxygen;  in  the  oxide  of  zinc,  32.5  grams  of  zinc  are  united  to 
8  grams  of  oxygen;  in  oxide  of  aluminum,  9  grams  of  aluminum 
are  united  to  8  grams  of  oxygen;  in  oxide  of  magnesium,  12 
grams  of  magnesium  are  united  to  8  grams  of  oxygen.  In  the 
chlorides  of  hydrogen,  sodium,  zinc,  aluminum,  and  magnesium, 
these  same  equivalents,  namely,  1,  23,  32.5,  9,  and  12,  respec- 
tively, are  each  united  with  35.45  parts  by  weight  of  chlorine. 
Thus,  one  might  look  upon  hydrochloric  acid  as  water  in  which 
the  8  parts  of  oxygen  that  are  united  to  1  part  of  hydrogen  are 
replaced  by  35.45  parts  of  chlorine.  Similarly,  for  instance, 
oxide  of  magnesium  might  be  regarded  as  water  in  which  each 
gram  of  hydrogen  is  replaced  by  12  grams  of  magnesium;  or 
the  chloride  of  magnesium  might  be  looked  upon  as  hydrochloric 
acid  in  which  each  gram  of  hydrogen  is  replaced  by  12  grams 
of  magnesium;  sodium  chloride  might  be  viewed  as  sodium 


FUNDAMENTAL  LAWS   AND   THEORIES  61 

oxide  in  which  every  8  grams  of  oxygen  are  replaced  by  35.45 
grams  of  chlorine,  etc. 

It  thus  becomes  clear  that  these  so-called  chemical  equivalents 
are  the  weights  of  the  elements  that  enter  into  chemical  com- 
bination with  one  another.  So  for  each  chemical  element  a 
definite  weight  can  be  found  experimentally,  in  which,  or  in 
multiples  of  which,  it  enters  into  combination  with  other  ele- 
ments. A  table  of  such  weights  is  called  a  table  of  combining 
weights  or  chemical  equivalents.  While  the  actual  numerical 
value  of  such  combining  weights  or  chemical  equivalents  will 
vary  according  to  the  unit  of  comparison  chosen,  the  relative 
values  of  the  combining  weights  are  fixed  and  invariable.  So 
we  might  arbitrarily  agree  to  call  the  combining  weight  of  oxy- 
gen 100,  and  then  determine  the  combining  weight  of  the  other 
elements  accordingly.  Thus  in  water  100  parts  of  oxygen  are 
combined  with  12.5  parts  of  hydrogen,  and  the  latter  figure 
would  be  the  combining  weight  of  hydrogen,  if  that  of  oxygen 
is  taken  as  100.  Similarly,  on  this  basis,  the  combining  weight 
of  chlorine  would  be  443.1,  for  in  hydrochloric  acid  12.5  parts 
of  hydrogen  by  weight  are  combined  with  443.1  parts  of  chlo- 
rine. As  a  matter  of  fact,  the  great  Swedish  chemist  Berzelius, 
who  lived  in  the  former  half  of  the  nineteenth  century,  and  to 
whose  careful,  painstaking  labors  we  owe  the  first  really  reliable 
table  of  the  combining  weights  of  the  chemical  elements,  actually 
computed  the  values  he  found  experimentally  by  choosing  arbi- 
trarily the  combining  weight  of  oxygen  as  100.  On  the  other 
hand,  Wollaston,  an  English  chemist  and  contemporary  of  Ber- 
zelius, arranged  the  table  on  the  basis  of  oxygen  equals  10.  It 
is  clear  that  one  might  arrange  such  a  table  by  taking  the  com^ 
bin  ing  weight  of  any  element  equal  to  any  fixed  number,  and 
then  computing  the  combining  weights  of  the  other  elements 
accordingly,  from  experimental  results  gathered  by  ascertaining 
the  ratios  in  which  the  elements  combine  with  one  another  by 
weight.  At  the  beginning  of  the  nineteenth  century,  John  Dai- 
ton  suggested  that  the  table  of  combining  weights  be  arranged 
by  arbitrarily  choosing  the  combining  weight  of  hydrogen  equal 
to  one.  Since  hydrogen  is  the  lightest  of  substances,  and  the 
quantity  by  weight  in  which  it  enters  into  combination  with 
other  elements  is  the  smallest,  it  is  clear  that  if  its  combining 
weight  is  chosen  as  unity,  the  combining  weights  of  all  the 


62  OUTLINES   OF   CHEMISTRY 

other  elements  would  be  greater  than  unity.  Now,  as  we  have 
seen,  hydrogen  combines  directly  with  but  very  few  elements, 
whereas  oxygen  is  a  very  active  substance  chemically,  and  com- 
bines directly  with  most  of  the  other  elements.  For  this  reason 
a  large  number  of  the  combining  weights  were  determined  from 
the  compounds  with  oxygen;  and  this  fact  led  Berzelius  to  take 
that  element  as  the  standard  for  his  table,  fixing  the  value  of 
its  combining  weight  arbitrarily  at  100,  as  already  stated. 
However,  the  distinct  advantage  of  placing  the  combining 
weight  of  hydrogen  equal  to  unity  was  so  apparent,  that  for  a 
long  time  tables  were  arranged  on  this  basis.  In  computing 
the  values  of  a  table  of  combining  weights  on  the  basis  that  the 
combining  weight  of  hydrogen  equals  1,  the  ratio  by  weight  in 
which  oxygen  and  hydrogen  combine  clearly  plays  a  very  im- 
portant role;  for  since  most  of  the  elements  do  not  combine 
with  hydrogen  directly,  their  combining  weights  are  actually 
determined  from  the  ratios  in  which  they  unite  with  oxygen. 
Consequently  the  ratio  in  which  the  latter  element  unites  with 
hydrogen  is  used  to  compute  the  combining  weights  on  the 
basis  of  hydrogen  equals  unity.  Now  since  hydrogen  is  a  very 
light  gas  and  consequently  difficult  to  weigh  with  accuracy  in 
large  quantities,  and  moreover  rather  difficult  to  obtain  in  a 
high  state  of  purity,  the  ratio  in  which  oxygen  and  hydrogen 
unite  to  form  water  has  been  the  object  of  a  great  deal  of  re- 
search by  a  goodly  number  of  able,  painstaking  chemists.  And 
every  time  this  ratio  was  determined  to  a  greater  degree  of 
accuracy,  nearly  all  the  values  in  the  table  of  combining  weights 
had  to  be  recomputed.  To  avoid  this,  most  chemists  have  of 
recent  years  returned  to  the  oxygen  basis,  choosing  the  combin- 
ing weight  of  oxygen  arbitrarily  as  equal  to  8.  On  this  basis, 
the  combining  weight  of  hydrogen  is  very  nearly  1.008,  accord- 
ing to  the  latest  researches.  By  this  procedure,  we  retain  the 
practical  advantage  of  having  hydrogen  as  the  unit  arid  avoid 
the  necessity  of  recalculating  nearly  the  whole  of  the  table  every 
time  the  ratio  in  which  oxygen  and  hydrogen  unite  is  redeter- 
mined  with  greater  accuracy.  Such  experimental  redetermina- 
tion  will  then  only  necessitate  a  recomputation  of  the  combining 
weight  of  hydrogen.  From  what  has  been  said,  it  is  evident, 
that  if  the  combining  weight  of  hydrogen  is  arbitrarily  taken 
as  equal  to  1,  that  of  oxygen  is  nearly  7.94. 


FUNDAMENTAL  LAWS  AND   THEORIES  63 

The  following  is  a  table  of  the  combining  weights  or  chemical 
equivalents  which  have  mainly  engaged  our  attention  thus  far 
in  this  book,  computed  on  the  basis  of  oxygen  equals  8  :  — 


Hydrogen 1.008 

Oxygen 8.00 

Chlorine 35.46 

Sulphur 16.03 

Sodium                              .  23.00 


Magnesium 12.16 

Aluminum 9.03 

Zinc 32.68 

Iron 27.92 

Lead  103.55 


For  purposes  of  discussion,  figures  rounded  off  to  the  first 
decimal  place  are  commonly  used. 

Chemical  Symbols.  —  Thus  far  in  all  our  considerations  we 
have  written  the  names  of  all  elements  and  compounds.  This  is 
a  somewhat  cumbersome  procedure,  and  chemists  have  adopted  a 
system  of  symbols  for  designating  substances.  Even  long  before 
the  rise  of  modern  chemistry,  which  dates  from  the  work^of  La- 
voisier at  the  beginning  of  the  nineteenth  century,  chemists  had 
been  wont  to  designate  substances  by  means  of  certain  symbols. 
So  the  moon  D  was  the  symbol  of  silver,  v  was  the  symbol  for 
water,  G.was  the  symbol  for  a  salt,  etc.  Such  arbitrary  signs 
were  still  used  by  the  great  Swedish  chemist  Scheele,  who  was 
a  contemporary  of  George  Washington.  It  was  Berzelius  who, 
in  the  early  part  of  the  nineteenth  century,  made  the  sugges- 
tion that  each  of  the  chemical  elements  be  designated  by  the 
first  letter  of  the  name  of  the  element,  and  that  whenever  the 
names  of  two  or  more  elements  begin  with  the  same  letter,  a 
second  characteristic  letter  from  the  name  of  the  element  be 
added  to  the  symbol.  This  is  still  the  usage  at  the  present 
time.  Thus,  the  symbol  of  hydrogen  is  H,  that  of  oxygen  O, 
that  of  carbon  C,  that  of  chlorine  Cl,  that  of  cobalt  Co,  etc.  In 
some  cases  the  symbols  are  derived  from  the  Latin  names  of 
the  elements  in  the  manner  mentioned.  So  the  symbol  for 
silver,  argentum,  is  Ag ;  for  sodium,  natrium,  Na ;  for  mer- 
cury, hydrargyrum,  Hg  ;  for  lead,  plumbum,  Pb  ;  etc.  A  com- 
plete list  of  all  the  names  and  symbols  of  the  elements  will  be 
given  later. 

Having  thus  adopted  the  symbols  for  the  elements,  the  com- 
pounds would  naturally  be  designated  by  simply  writing  side 
by  side  the  symbols  of  the  elements  that  occur  in  the  com- 
pounds. But  since  the  elements  always  enter  into  combina- 
tion in  the  ratio  of  their  combining  weights,  it  is  easy  to  have 


64  OUTLINES   OF  CHEMISTRY 

the  symbol  of  a  compound  stand  for  both  its  qualitative  and  its 
quantitative  composition.  This  is  readily  accomplished  by 
letting  the  symbol  of  each  element  stand  for  not  only  the 
verbal  name  of  that  element,  but  also  for  its  combining  weight. 
Thus  the  symbols  H  and  Cl  would  stand  not  only  for  hydro- 
gen and  chlorine,  but  also  for  1  part  of  hydrogen  by  weight, 
and  35.45  parts  of  chlorine  by  weight,  respectively.  And  so 
the  symbol  HC1  stands  not  only  for  hydrochloric  acid,  but  it 
also  tells  us  that  in  that  compound  1  part  of  hydrogen  is  com- 
bined with  35.45  parts  of  chlorine  by  weight.  Likewise,  the 
symbol  for  common  salt,  NaCl,  denotes  that  this  compound  con- 
sists of  sodium  and  chlorine  united  in  the  proportions  of  23 
parts  of  sodium  to  35.45  parts  of  chlorine  by  weight.  Again, 
the  symbol  HO,  which  was  formerly  used  for  water,  denoted 
that  this  compound  consists  of  hydrogen  and  oxygen  united  in 
the  proportions  of  1  to  8  by  weight.  The  symbol  of  ferrous 
sulphide,  FeS,  denoted  that  in  this  compound  iron  and  sulphur 
are  present  and  in  the  proportion  of  28  parts  of  iron  to  16  parts 
of  sulphur  by  weight.  The  symbol  FeS2,  the  2  being  used  as 
a  subscript  to  the  S,  was  the  symbol  for  pyrite,  and  denoted 
that  in  it  28  parts  of  iron  are  combined  with  2  times  16  parts 
of  sulphur,  i.e.  two  combining  weights  of  sulphur.  In  general, 
whenever  more  than  one  combining  weight  of  an  element  enters 
into  the  compound,  that  fact  is  indicated  by  the  appropriate 
figure  used  as  a  subscript.  So,  for  instance,  the  symbol  for 
red  lead  is  Pb3O4,  indicating  that  in  that  compound  3  combin- 
ing weights  (3  x  103.5)  of  lead  are  united  with  4  combining 
weights  (4  x  8)  of  oxygen.  This  mode  of  designating  chemical 
compounds  by  having  the  symbols  stand  for  equivalent  or  com- 
bining weights  was  in  vogue  for  many  years  ;  and  with  a  slight 
modification  it  is  still  in  use  at  present.  The  nature  of  this  modi- 
fication lies  merely  in  the  fact  that  we  do  not  always  designate 
the  chemical  equivalent  by  the  symbol  of  the  element,  but  fre- 
quently the  symbol  stands  for  some  other  simple  multiple  of  the 
chemical  equivalent,  for  reasons  that  will  presently  be  explained. 
The  Atomic  Theory  of  Matter.  —  The  fact  that  the  chemical 
elements  always  unite  in  definite  proportions  by  weight  in 
accordance  with  the  three  laws  of  definite,  multiple,  and  recip- 
rocal proportions,  finds  a  ready  explanation  in  a  simple  as- 
sumption as  to  the  nature  of  matter.  If  we  assume  that  each 


FUNDAMENTAL   LAWS   AND   THEORIES  65 

elementary  substance  is  made  up  of  extremely  minute,  ultra- 
microscopic,  indivisible  particles,  termed  atoms  (from  the  Greek 
meaning  indivisible),  and  that  these  atoms  are  of  exactly  the 
same  weight  and  also  otherwise  alike  in  the  case  of  any  one 
element,  but  different  in  weight  and  other  properties  in  the 
case  of  different  elements,  and  that  chemical  compounds  are 
formed  by  the  union  of  the  atoms  of  the  various  elements  with 
one  another,  the  experimental  facts  expressed  in  the  laws  of 
definite,  multiple,  and  reciprocal  proportions  are  readily  ex- 
plained. So,  for  instance,  by  this  hypothesis,  hydrogen  would 
be  considered  as  made  up  of  minute  particles,  atoms,  all  of  the 
same  weight  and  otherwise  also  alike.  Chlorine  would  similarly 
be  regarded  as  composed  of  atoms  which  are  of  the  same  weight 
and  otherwise  alike  among  themselves,  but  quite  different  in 
weight  and  other  properties  from  the  atoms  of  hydrogen  or 
those  of  any  other  element.  Each  element  would  similarly  be 
composed  of  atoms  that  are  alike  in  weight  and  otherwise,  but 
different  in  weight  and  other  characteristics  from  the  atoms  of 
all  other  elements. 

Since  the  atoms  of  each  element  are  assumed  to  be  indivisible, 
in  forming  compounds  a  whole  number  of  atoms  of  one  element 
must  always  unite  with  a  whole  number  of  atoms  of  another 
element  or  elements.  Consequently  the  proportions  by  weight 
in  which,  for  instance,  two  elements  A  and  B  unite  with  each 
other  to  form  a  compound  AB,  are  proportional  to  the  atomic 
weights  of  A  and  B  ;  in  other  words,  the  combining  weights  of  the 
elements  are  proportional  to  the  atomic  weights.  So,  for  example, 
1  gram  of  hydrogen  unites  with  35. 45"  grams  of  chlorine  to 
form  36.45  grams  of  hydrochloric  acid,  consequently,  in  the 
light  of  the  atomic  theory,  1  gram  of  hydrogen  must  contain 
a  definite  whole  number  of  atoms  of  hydrogen,  and  similarly 
the  35.45  grams  of  chlorine  must  contain  a  definite  whole  num- 
ber of  atoms  of  chlorine.  If  we  let  x  represent  the  actual 
weight  of  one  of  the  hydrogen  atoms,  and  n  the  number  of 
atoms  of  hydrogen  in  1  gram  of  hydrogen,  then  xn  equals  1 
gram.  Similarly,  if  we  let  y  represent  the  weight  of  1  atom 
of  chlorine  and  n'  the  number  of  atoms  of  chlorine  in  35.45 
grams  of  that  gas,  then  yn1  equals  35.45  grams.  We  may 
consequently  write  the  relation  — 

xn  :  ynf :  :  1  gram  :  35.45  grams. 


66  OUTLINES   OF  CHEMISTRY 

It  is  clear  that  we  might  deduce  a  similar  equation  in  the  case 
of  the  union  of  any  two  or  more  of  the  chemical  elements.  Let 
us  now  view  this  equation  more  closely.  It  contains  four  un- 
known quantities  :  namely,  the  atomic  weight  of  hydrogen  #,  the 
atomic  weight  of  chlorine  ?/,  the  number  of  atoms  of  hydrogen  n, 
in  1  gram  of  Irydrogen,  and  the  number  of  atoms  of  chlorine  nr, 
in  35.45  grams  of  chlorine.  Of  course,  neither  of  these  values 
can  be  ascertained  from  the  equation  as  it  stands.  If  now  we 
arbitrarily  choose  some  definite  value  for  either  x  or  y,  say  we 
assume  with  Dalton  the  atomic  weight  of  hydrogen  as  1,  the  x 
will  disappear  from  our  equation  as  an  unknown  quantity  ;  still 
we  should  have  the  three  unknown  quantities  n,  n1,  and  y 
present  in  the  equation.  There  is  no  way  of  determining  how 
many  atoms  of  hydrogen  unite  with  how  many  atoms  of  chlo- 
rine in  forming  hydrochloric  acid,  and  so  it  is  customary  to 
make  the  simplest  possible  assumption  here,  namely,  that  1 
atom  of  hydrogen  unites  with  1  atom  of  chlorine  in  forming 
a  particle  of  hydrochloric  acid.  On  the  basis  of  this  assump- 
tion, it  becomes  clear  that  1  gram  of  hydrogen  would  contain 
as  many  atoms  of  hydrogen  as  35.45  grams  of  chlorine  contain 
atoms  of  chlorine ;  or  in  other  words,  in  our  equation  n  equals 
nr.  Since  we  have  assumed  that  x  equals  1,  and  also  that  n 
equals  n\  the  equation  becomes  — 

1 :  y :  :  1  gram  :  35.45  grams ;  whence  y  equals  35.45. 

That  is,  the  atomic  weight  of  chlorine  is  35.45,  if  the  atomic 
weight  of  hydrogen  is  assumed  to  be  1,  and  it  is  further  assumed 
that  in  forming  hydrochloric  acid  1  atom  of  hydrogen  unites 
with  1  atom  of  chlorine.  Similarly,  the  proportions  by  weight 
in  which  hydrogen  and  oxygen  unite  to  form  water,  namely, 
1  to  8  (if  we  were  to  make  the  simplest  assumption,  as  Dalton 
did,  that  in  forming  water  1  atom  of  hydrogen  unites  with  1 
atom  of  oxygen),  lead  to  the  conclusion  that  the  atomic  weight 
of  oxygen  is  8.  And  this  value  was  assigned  to  it  by  Dalton, 
though  it  is  not  the  one  used  at  present,  as  will  be  explained 
shortly.  However,  if  we  thus  take  the  atomic  weight  of 
hydrogen  as  1,  that  of  oxygen  as  8,  and  that  of  chlorine  as 
35.45,  then  since  in  chlorine  monoxide  every  35.45  grams  of 
chlorine  are  combined  with  8  grams  of  oxygen,  we  should  have 
in  this  compound  1  atom  of  chlorine  united  with  1  atom  of 


FUNDAMENTAL  LAWS   AND   THEORIES  67 

oxygen.  In  the  chlorine  peroxide  again,  in  which  every  35.45 
grams  of  chlorine  are  combined  with  32  grams  of  oxygen,  we 
should  have  1  atom  of  chlorine  combined  with  4  atoms  of  oxygen. 
Thus  on  this  basis  the  formulae  for  hydrochloric  acid,  water, 
chlorine  monoxide,  and  chlorine  peroxide  would  be  HC1,  HO, 
CIO,  and  C1O4,  respectively.  In  these  formulae  the  symbols 
of  the  elements  simply  stand  for  the  combining  or  equivalent 
weights.  Thus  by  introducing  the  assumptions :  (1)  that  the 
elements  are  made  up  of  atoms,  (2)  that  the  atomic  weight  of 
hydrogen  is  1,  (3)  that  in  water  1  atom  of  hydrogen  is  united 
with  1  atom  of  oxygen,  and  (4)  that  in  hydrochloric  acid  1  atom 
of  chlorine  is  united  with  1  atom  of  hydrogen,  we  simply  arrive 
at  the  conclusion  that  the  combining  weights  are  the  relative 
atomic  weights ;  the  lowest  combining  weight  of  an  element 
found  in  any  compound  into  which  it  enters  being,  of  course, 
taken  as  the  atomic  weight.  This  system  of  using  the  equiva- 
lent weights  as  the  atomic  weights  was  employed  by  many 
chemists  during  the  first  half  of  the  19th  century. 

Thus,  in  the  atomic  theory,  the  law  of  definite  proportions 
finds  a  ready  explanation,  for  each  compound  would  always 
contain  the  same  relative  number  of  atoms  of  each  of  the  ele- 
ments of  which  it  is  composed.  The  law  of  multiple  propor- 
tions is  readily  explained  by  the  theory,  since  according  to  it  a 
fixed  number  of  atoms  of  one  element  can  only  combine  with 
one  atom  or  some  other  whole  number  of  atoms  of  another  ele- 
ment. And  finally,  the  law  of  reciprocal  proportions  also  is 
easily  accounted  for  by  the  theory,  since  according  to  it  the 
weight  of  an  atom  of  any  one  element  is  constant  and  different 
from  that  of  any  other  element,  and  combination  can  only  take 
place  by  whole  numbers  of  atoms,  from  which  it  follows  that 
the  proportion  by  weight  in  which  an  element  occurs  in  one 
compound  will  be  either  the  same  as,  or  some  multiple  of,  the 
proportion  in  which  it  occurs  in  any  other  compound. 

It  should  thus  be  borne  in  mind  that  the  modern  atomic 
theory  of  matter  is  based  upon  the  weight  relations  that  obtain 
when  the  elements  unite  chemically,  and  these  weight  relations 
are  expressed  in  the  three  fundamental  stoichiometrical  laws. 

It  is  of  interest  to  note  that,  though  Dalton  promulgated  the 
modern  atomic  theory  in  1802,  the  basis  for  that  theory  has 
been  furnished  by  chemists  of  three  different  nations ;  for  the 


68  OUTLINES   OF   CHEMISTRX 

law  of  definite  proportions  was  discovered  by  Lavoisier  and 
Proust,  the  law  of  multiple  proportions  by  Dalton,  and  the  law 
of  reciprocal  proportions  by  Richter. 

The  atomic  conception  of  matter  was  not  original  with 
Dalton ;  indeed,  its  origin  dates  back  to  the  times  of  classical 
Greece.  Democritus,  Epicurus,  and  Leucippus  taught  that 
matter  is  made  up  of  indivisible  particles  or  atoms,  whereas 
according  to  the  doctrine  of  Anaxagoras,  matter  is  infinitely 
divisible.  However,  the  atomic  conception  of  matter  of  the 
Greeks  was  a  mere  metaphysical  speculation,  not  founded  upon 
experimental  facts.  It  was  Dalton  who  first  used  the  concep- 
tion of  the  atomic  nature  of  matter  in  explaining  actual  facts 
established  by  experiments,  and  to  him  consequently  we  rightly 
ascribe  the  origin  of  the  modern  atomic  theory. 

Difference  between  Theory  and  Law.  —  The  student  must 
always  clearly  bear  in  mind  the  distinction  between  a  theory 
on  the  one  hand,  and  facts  and  laws  on  the  other  hand.  Facts 
are  the  results  of  actual  observation  and  experiment.  When  a 
large  number  of  similar  facts  have  been  found  and  these  are  ex- 
pressed in  a  general  statement,  the  latter  is  a  law.  Thus  we 
actually  find  that  the  composition  of  salt,  water,  lime,  sal  am- 
moniac, etc.,  is  constant,  no  matter  when  or  where  prepared. 
We  have  here  a  series  of  facts.  If  now  we  formulate  this  into 
the  general  statement,  that  a  chemical  compound  always  has 
the  same  composition,  we  have  a  law.  A.  theory  or  hypothesis, 
however,  is  neither  a  fact  nor  a  general  statement  of  fact,  it  is 
merely  an  assumption  made  for  the  purpose  of  correlating,  explain- 
ing, or  accounting  for  facts  that  have  been  collected  and  formulated 
into  laws.  So  the  atomic  theory  is  a  theory  which  enables  us 
to  correlate  and  comprehend  better  the  facts  expressed  in  the 
stoichiometrical  laws.  A  theory,  however,  not  only  enables  one 
to  see  facts  in  their  relations  and  thus  satisfies  the  natural  craving 
of  the  human  intellect  for  a  better  comprehension  of  things  observed, 
but  it  also  suggests  new  avenues  of  inquiry  and  experimentation 
by  means  of  which  further  facts  may  be  acquired.  A  theory  is 
thus  a  powerful  stimulus  to  scientific  research,  and  is  conse- 
quently of  almost  inestimable  value.  On  the  other  hand,  it  is 
to  be  remarked  that  theories  by  implication  also  suggest  that  it  is 
useless  for  actual  inquiry  to  proceed  in  certain  directions,  and  that 
certain  things  are  impossible,  when  after  all  they  are  quite  possible, 


FUNDAMENTAL   LAWS   AND   THEORIES  69 

and  thus  a  theory  may  be  a  bar  to  progress.  Carefully  ascertained 
facts  formulated  into  laws  constitute  the  unchangeable,  the 
eternal  part  of  science.  Theories  and  hypotheses  on  the  other 
hand  are  the  changeable,  the  ephemeral  part  of  science  ;  for  the 
views  we  entertain  concerning  the  relationship  of  natural  phe- 
nomena frequently  change  as  new  facts  are  acquired.  A  theory 
is  a  cord  by  which  the  precious  pearls  of  truth  are  held  together, 
but  when  the  pearls  found  become  too  numerous  or  too  heavy 
so  that  the  old  cord  can  no  longer  hold  them  together,  it  must 
be  discarded,  and  the  pearls  must  ultimately  be  arranged  on  a 
new  cord  of  adequate  length  and  strength.  Thus  theories  and 
hypotheses  are  frequently  discarded  in  science.  In  fact,  the 
pathway  of  the  progress  of  science  is  strewn  with  defunct 
theories.  As  we  continue  our  considerations,  we  shall  see  that 
the  atomic  theory,  simple  and  even  crude  as  it  seems,  has  been 
in  a  high  degree  useful  in  correlating  even  facts  other  than 
those  upon  which  it  is  actually  based,  and  has  suggested  many 
new  avenues  of  further  experimental  inquiry.  It  has  thus  ful- 
filled in  a  high  degree  the  function  of  a  theory. 

The  Law  of  Combination  of  Gases  by  Volume.  —  It  will  be  re- 
called that  Gay-Lussac  discovered  the  law  that  when  gases  com- 
bine to  form  chemical  compounds  the  volumes  of  the  gases  that 
enter  into  combination  bear  a  simple  relation  to  one  another ;  and 
if  the  product  formed  be  gaseous,  its  volume  also  bears  a  simple 
relation  to  the  volumes  of  the  original  gases.  This  law  was  estab- 
lished at  about  the  time  when  Dalton  formulated  the  atomic 
theory.  In  viewing  the  fact  that  1  volume  of  hydrogen  and 
1  volume  of  chlorine  unite  to  form  hydrochloric  acid  gas,  in 
the  light  of  the  atomic  theory  of  Dalton,  according  to  which 
hydrogen  and  chlorine  are  made  up  of  atoms  and  1  atom  of 
the  one  unites  with  1  atom  of  the  other  to  form  one  particle 
of  hydrochloric  acid,  it  follows  at  orice  that  1  volume  of 
hydrogen  must  contain  exactly  as  many  atoms  of  hydrogen  as 
the  same  volume  of  chlorine  contains  atoms  of  the  latter  ele- 
ment ;  for  were  this  not  the  case,  there  would  be  either  hydro- 
gen or  chlorine  left  uncombined  when  exactly  equal  volumes 
act  on  each  other  chemically.  Thus  the  idea  was  naturally 
formed  that  equal  volumes  of  gases  under  the  same  conditions 
of  temperature  and  pressure  contain  the  same  number  of  atoms. 
This  is,  of  course,  not  a  law,  but  simply  an  hypothesis  evolved 


70  OUTLINES   OF   CHEMISTRY 

to  explain  the  law  of  Gay-Lussac  of  combination  of  gases  by 
volume.  To  the  hypothesis  in  the  form  stated,  Berzelius  inter- 
posed a  serious  objection.  Thus  he  called  attention  to  the  fact 
that  when  1  volume  of  hydrogen  and  1  volume  of  chlorine  unite, 
2  volumes  of  hydrochloric  acid  are  formed ;  and  if  1  atom  of 
hydrogen  unites  with  1  atom  of  chlorine,  there  will  of  course 
be  formed  as  many  particles  of  hydrochloric  acid  as  there  are 
particles  of  hydrogen,  or  what  amounts  to  the  same  thing,  as 
there  are  particles  of  chlorine.  Therefore,  if  equal  volumes 
of  hydrogen,  chlorine,  and  hydrochloric  acid  contain  the  same 
number  of  atoms  or  particles,  the  hydrochloric  acid  formed  by 
the  union  of  equal  volumes  of  hydrogen  and  chlorine  ought 
to  occupy  the  same  volume  as  the  original  hydrogen ;  that  is, 
it  ought  to  occupy  one  half  of  the  volume  that  it  actually  does 
occupy.  To  hold  the  volume  hypothesis,  a  scheme  consequently 
had  to  be  proposed  whereby  1  atom  of  hydrogen  would  unite 
with  1  atom  of  chlorine  and  form  2  particles  of  hydrochloric 
acid;  for  only  then  the  actual  volume  relations  that  obtain 
when  the  latter  substance  is  formed  by  the  union  of  hydrogen 
and  chlorine  would  be  accounted  for.  Such  a  scheme  was  pro- 
posed in  1811  by  Amadeo  Avogadro,  who  was  then  professor 
of  physics  at  the  University  of  Turin.  He  made  the  bold 
assumption  that  the  particles  of  hydrogen  gas  really  are  double 
atoms,  that  is,  that  they  consist  of  2  atoms  joined  together, 
and  that  the  particles  of  chlorine  gas  are  similarly  made  up 
each  of  2  chlorine  atoms.  These  double  atoms  of  hydrogen 
and  double  atoms  of  chlorine  he  called  molecules,  and  then  stated 
the  hypothesis  as  follows  :  Equal  volumes  of  all  gases  under  the 
same  conditions  of  temperature  and  pressure  contain  the  same 
number  of  molecules.  This  is  known  as  Avogadro's  hypothesis. 
It  is  very  important  in  chemistry.  Thus,  Avogadro  considered 
the  molecule  of  hydrogen  as  H2,  and  the  molecule  of  chlorine 
as  C12;  and  when  these  react  with  each  other  we  should  have  — 

H2        +         C12        =2HC1. 

1  molecule  -f- 1  molecule  =  2  molecules. 
1  volume    +  1  volume     =  2  volumes. 

On  this  basis,  there  would  be  twice  as  many  molecules  of 
hydrochloric  acid  formed  as  there  were  molecules  of  hydrogen 
or  molecules  of  chlorine,  and  consequently  one  would  expect 


FUNDAMENTAL   LAWS   AND   THEORIES  71 

the  hydrochloric  acid  formed  to  occupy  twice  the  volume  of  the 
original  hydrogen;  or,  what  is  the  same,  twice  the  volume  of 
the  chlorine. 

Let  us  now  review  the  volume  relations  that  obtain  when 
hydrogen  and  oxygen  combine  to  form  water  vapor.  We  have 
by  experiment  — 

2  volumes  of  hydrogen  -f-  1  volume  of  oxygen 

=  2  volumes  of  water  vapor. 

If,  now,  we  desire  to  hold  Avogadro's  hypothesis,  we  clearly 
must  assume  the  molecules  of  hydrogen,  oxygen,  and  water  so 
constituted  that :  — 

2  molecules  of  hydrogen  +  1  molecule  of  oxygen 

=  2  molecules  of  water  vapor. 

But  in  connection  with  the  synthesis  of  hydrochloric  acid,  it 
was  already  assumed  that  the  molecule  of  hydrogen  consists  of 
2  atoms  (i.e.  that  it  is  H2) ;  we  must  consequently  adhere  consist- 
ently to  this  assumption  wherever  hydrogen  gas  is  considered. 
Now  if  we  assume  that  the  oxygen  molecule  is  made  up  of  2 
atoms  of  oxygen  (i.e.  is  O2),  the  volume  relations  in  the  case  of 
the  synthesis  of  water  are  readily  explained ;  the  water  mole- 
cule then,  however,  must  be  considered  as  composed  of  2  atoms 
of  hydrogen  and  1  atom  of  oxygen.  In  form  of  an  equation 
we  should  have  — 

2  volumes  of  hydrogen  +  1  volume  of  oxygen  =  2  volumes  of  water  vapor. 
2  molecules  of  hydrogen  +  1  molecule  of  oxygen  =  2  molecules  of  water  vapor. 
i.e.  2H2  +  O2  =2H2O. 

Thus  it  appears  that  if  we  hold  Avogadro's  hypothesis,  and 
assume  with  him,  that  in  hydrogen  gas  and  oxygen  gas  there 
are  molecules  that  consist  of  2  atoms  of  these  respective  ele- 
ments, we  are  bound  to  conclude  that  the  molecule  of  the  very 
common  compound  water  is  not  made  up  of  1  atom  of  hydro- 
gen united  with  1  atom  of  oxygen,  but  rather  of  2  atoms  of 
hydrogen  united  with  1  atom  of  oxygen.  We  should  thus  have 
to  assign  to  water  the  formula  H2O  instead  of  HO.  Further, 
since  by  weight  1  part  of  hydrogen  unites  with  8  parts  of  oxy- 
gen, and  since  we  assume  with  Dalton  that  the  atomic  weight 
of  hydrogen  is  1,  we  consequently  must  assume  the  atomic 
weight  of  oxygen  as  equal  to  16  instead  of  8. 


72  OUTLINES   OF   CHEMISTRY 

All  this  seemed  to  many  chemists  of  the  early  part  of  the 
nineteenth  century  as  a  set  of  rather  violent  changes  to  make. 
In  fact,  Avogadro's  molecules,  his  double  atoms  as  they  were 
frequently  termed  in  the  literature,  were  not  regarded  seriously 
by  many  able  chemists  for  nearly  half  a  century ;  they  simply 
continued  to  work  with  the  tables  of  equivalent  or  combining 
weights  in  which  the  value  for  oxygen  was  8.  One  of  the 
reasons  why  Avogadro's  hypothesis  was  laid  aside  for  a  time 
was  that  in  the  study  of  the  ammonium  salts  certain  apparent 
contradictions  were  met.  These  will  be  considered  when  those 
salts  are  discussed.  Suffice  it  here  to  say  that  Avogadro's 
hypothesis  has  gained  general  acceptance  and  is  now  commonly 
regarded  as  of  vital  consequence  in  chemistry.  Thus,  while  the 
atomic  theory  is  based  upon  the  weight  relations  that  obtain  when 
substances  unite  chemically,  the  molecular  theory  came  into  being 
as  a  consequence  of  the  acceptance  of  Avogadro's  hypothesis,  which 
in  turn  grew  out  of  G-ay-Lussac's  law  of  the  combination  of  gases 
by  volume.  Avogadro's  hypothesis  is  further  supported  by  the 
fact  that  all  gases  contract  and  expand  alike  under  the  same 
changes  of  temperature  and  pressure. 

Molecular  Weight  Determinations. — If  equal  volumes  of  all 
gases  under  the  same  conditions  of  temperature  and  pressure 
contain  the  same  number  of  molecules,  it  is  clear  that  the 
weights  of  equal  volumes  of  gases  are  to  one  another  as  the  mo- 
lecular weights  of  the  gases.  To  fully  appreciate  Avogadro's 
hypothesis,  the  student  must  bear  in  mind  that  on  the  basis  of 
the  molecular  theory  a  gas  consists  of  molecules  that  are  en- 
tirely remote  from  one  another.  So,  for  example,  if  the  mole- 
cules of  a  gas  were,  say  by  pressure,  all  crowded  together  so 
that  they  touched  one  another,  they  would  occupy  but  a  small 
portion  of  the  volume  originally  occupied  by  the  gas.  In  other 
words,  the  actual  volume  of  a  gas  consists  largely  of  space  or 
interstices  between  the  molecules,  as  it  were.  Bearing  this  in 
mind,  it  is  clear  that  equal  volumes  of  gases  may  well  contain 
equal  numbers  of  molecules,  though  the  individual  molecules 
of  each  of  the  gases  may  occupy  very  different  volumes.  Ac- 
cording to  Avogadro's  hypothesis,  the  volume  of  any  gas,  at 
constant  temperature  and  pressure,  depends  not  upon  the  size, 
weight,  or  kind  of  molecules  it  contains,  but  solely  upon  the 
number  of  molecules  present.  Hence  at  constant  temperature  and 


FUNDAMENTAL   LAWS  AND  THEORIES  73 

pressure,  the  volumes  of  any  two  gases  are  to  each  other  as  the 
number  of  molecules  in  the  volumes.  Thus,  for  instance,  10 
liters  of  any  gas  contain  10  times  as  many  molecules  as  1 
liter  of  the  same  gas  or  any  other  gas.  Again,  take  any  vol- 
ume of  hydrogen,  say  1  liter,  and  compare  its  weight  with  the 
weight  of  the  same  volume  of  some  other  gaseous  substance, 
say  chloroform  vapor,  at  the  same  temperature  and  pressure. 
Now  since  by  Avogadro's  hypothesis  each  of  these  volumes 
contains  the  same  number  of  molecules,  which  we  shall  call  n, 
then  we  should  have  — 
The  wt.  of  1  liter  of  Chloroform  Vapor  :  wt.  of  1  liter  Hydrogen  :  :  ran  :  2  w, 

where  m  is  the  molecular  weight  of  chloroform,  and  2  that  of 
hydrogen.  From  the  equation, 

0      wt.  of  liter  of  Chloroform  Vapor 

/vyi   — —    £    X    ~ " * 

1  wt.  of  liter  of  Hydrogen 

That  is,  the  molecular  weight  of  chloroform  is  equal  to  twice  its 
vapor  density  as  compared  with  hydrogen.  Hence  the  rule  for 
finding  molecular  weight  of  any  gas :  find  the  density  of  the 
gas  with  respect  to  hydrogen,  and  multiply  the  result  by  2. 
The  volume  occupied  by  2  grams  of  hydrogen  under  standard 
conditions,  0°  and  760  mm.  pressure,  is  22.38  liters.  The  same 
volume  is  occupied  by  32  grams  of  oxygen,  70.9  grams  of 
chlorine,  18  grams  of  water  vapor,  and  in  short,  by  the  molec- 
ular weight  in  grams  of  any  gaseous  substance  whatever,  under 
standard  conditions. ,  Consequently  22.38  liters  is  termed  the 
molecular  volume  of  all  gases.  We  may  also  state  that  to  find 
the  molecular  weight  of  any  gas,  we  simply  need  to  determine 
the  weight  of  22.38  liters  of  that  gas  at  0°  and  760  mm. 
pressure,  and  the  result  is  the  molecular  weight  in  grams. 
The  student  should  assure  himself  that  this  really  comes  to 
the  same  thing  as  finding  the  density  of  the  gas  with  respect  to 
hydrogen  and  multiplying  the  result  by  2. 

The  molecular  weights  of  substances  that  can  be  obtained  in 
the  vapor  state  can  thus  readily  be  determined.  But  there  are 
liquid  and  solid  substances  that  cannot  be  vaporized  without 
decomposition,  and  so  their  vapor  densities  cannot  be  deter- 
miined.  In  the  case  of  substances  that  can  be  dissolved,  it  is 
possible  to  make  molecular  weight  determinations  by  studying 
the  freezing  point,  boiling  point,  or  vapor  pressure  of  the  solu- 


74  OUTLINES   OF   CHEMISTRY 

tion.  This  will  be  explained  later  in  connection  with  the 
subject  of  solutions. 

Determination  of  Atomic  Weights.—  The  proportions  by  weight 
in  which  the  elements  combine  with  one  another  are  determined 
by  very  exact  chemical  analyses  of  the  compounds  containing 
the  elements  in  question ;  or,  if  the  latter  will  unite  directly, 
by  ascertaining  the  weights  of  the  elements  that  enter  into 
combination,  when  compounds  are  thus  formed  by  synthesis. 
The  values  so  found  are  the  combining  weights.  Dalton  took 
the  combining  weight  of  hydrogen  as  equal  to  unity,  and 
expressed  the  combining  weights  of  the  other  elements  on  this 
basis.  We  now,  for  reasons  already  stated  above,  take  the  com- 
bining weight  or  chemically  equivalent  weight  of  oxygen  as 
equal  to  8,  on  which  basis  that  of  hydrogen  equals  1.008. 

We  have  seen  that  Gay-Lussac's  law  of  combination  of  gases 
by  volume  led  to  Avogadro's  hypothesis, 'which  in  turn  led  to 
the  idea  of  molecules,  and  to  the  conception  that  a  molecule  of 
oxygen  consists  of  2  atoms.  This  further  made  it  necessary 
to  adopt  for  water  the  formula  H2O,  instead  of  HO,  and  con- 
sequently for  oxygen  the  atomic  weight  16  instead  of  8.  Thus 
the  experimental  fact  that  2  volumes  of  hydrogen  unite  with 

1  volume    of   oxygen   to    form    2    volumes    of   water    vapor 
really  determined  us  in  choosing  the  atomic  weight  of  oxygen 
as  16  instead  of  8 ;  in  other  words,  the  vapor  density  of  water 
has  really  been  used  in  deciding  whether  we  should  use  8  or 
some  multiple  of  that  figure  as  the  atomic  weight  of  oxygen. 
Similarly,  the  vapor  density  of  substances  has  in  many  other 
cases  been  used  in  choosing  the  atomic  weights,  the  combining 
weights  being  known.     Thus,  in  hydrochloric  acid  hydrogen 
and  chlorine  are  combined  in  the  proportions  of  1  to  35.45 
by  weight.     By  volume,  on  the  other  hand,  we  have,  1  volume 
of    hydrogen   uniting    with    1    volume    of    chlorine    to    form 

2  volumes   of   hydrochloric   acid.       We   have  seen   that   this 
volume   relation    led    Avogadro    to   distinguish    between   the 
atom  of  hydrogen,  H,  and  the  molecule  of  hydrogen,  H2,  and 
also  between  the  atom  of  chlorine,  Cl,  and  the  molecule  of 
chlorine,  C12 .     We  have  also  noted,  that  having  assumed  each 
of  the  molecules  of  hydrogen  and  chlorine  to  consist  of  2  atoms, 
the  composition  of  the  molecule  of  hydrochloric  acid  could  still 
be  expressed  by  the  simple  formula,  HC1 ;  and  thus  the  volume 


FUNDAMENTAL  LAWS   AND   THEORIES  75 

relations  that  obtain  when  hydrochloric  acid  is  formed  from  the 
elements  could  be  explained  in  the  light  of  Avogadro's  hypothe- 
sis, and  35.45  be  regarded  as  the  atomic  weight  of  chlorine. 
Here  again,  the  vapor  density  of  hydrochloric  acid  gas  has  been 
the  determining  factor  in  choosing  the  atomic  weight  of  chlorine 
as  35.45  rather  than  some  multiple  thereof. 

In  marsh  gas,  which  consists  of  hydrogen  and  carbon,  1  gram 
of  hydrogen  is  combined  with  every  3  grams  of  carbon.  The 
combining  weight  or  chemical  equivalent  of  carbon  is  therefore 
3.  How  now  proceed  to  ascertain  the  atomic  weight  of  carbon  ? 
Marsh  gas  is  8  times  heavier  than  hydrogen  ;  whence  22.38 
liters  of  marsh  gas  weigh  16  grams.  From  what  has  been 
stated  above,  16  is  therefore  the  molecular  weight  of  marsh 
gas.  But  in  16  grams  of  marsh  gas  there  are  4  grams  of 
hydrogen,  which  is  4  times  the  atomic  weight  of  hydrogen  in 
grams.  The  molecule  of  marsh  gas  therefore  contains  4  atoms 
of  hydrogen.  Now  if  in  the  case  of  this  compound,  which  of 
all  the  compounds  of  hydrogen  with  carbon  contains  the  least 
amount  of  carbon  by  weight  as  compared  with  the  hydrogen, 
we  assume  that  there  is  but  1  carbon  atom  to  the  4  hydrogen 
atoms,  we  .must  ascribe  to  the  carbon  atom  the  weight  of  12. 
Thus  it  is  clear  that  the  vapor  density  of  marsh  gas  has  not 
only  fixed  its  molecular  weight,  but  has  also  led  us  to  choose 
the  atomic  weight  of  carbon  as  12,  rather  than  as  some  other 
multiple  of  3,  the  combining  weight.  Further,  we  find  that 
by  thus  taking  12  as  the  atomic  weight  of  carbon,  the  com- 
position of  other  compounds  into  which  that  element  enters 
can  readily  be  expressed.  Of  compounds  of  carbon  with 
oxygen,  the  one  that  contains  the  least  carbon  as  compared 
with  the  oxygen  is  carbonic  acid  gas.  In  this  3  grams  of  car- 
bon are  combined  with  every  8  grams  of  oxygen.  The  gas  is 
22  times  heavier  than  hydrogen,  that  is  22.38  liters  of  carbonic 
acid  gas  weigh  44  grams,  and  the  molecular  weight  of  this  sub- 
stance is  consequently  44.  Since  it  contains  the  least  carbon 
of  any  of  the  known  compounds  of  oxygen  and  carbon,  it 
would  be  natural  to  hold  that  it  contains  but  1  atom  of  carbon. 
Now  if  the  atomic  weight  of  carbon  be  12,  that  of  oxygen  16, 
and  the  molecular  weight  of  carbonic  acid  gas  44,  we  have 
(since  in  44  grains  of  carbonic  acid  gas  there  are  12  grams  of 
carbon  and  32  grams  of  oxygen)  in  the  carbonic  acid  molecule 


76  OUTLINES   OF  CHEMISTRY 

1  atom  of  carbon  united  with  2  atoms  of  oxygen,  which  is  ex- 
pressed  by  the  formula  CO2.  Since  this  compound  contains 
but  two  elements,  it  is  a  so-called  binary  compound.  The 
names  of  all  binary  compounds,  i.e.  compounds  consisting  of 
but  two  elements,  end  in  ide.  Since  carbonic  acid  gas  contains 
twice  as  much  oxygen  as  the  lower  oxide  of  carbon,  carbon 
monoxide,  and  as  we  have  assigned  to  the  former  the  formula, 
CO2,  indicating  that  the  molecule  contains  2  atoms  of  oxygen, 
it  is  called  carbon  dioxide.  In  carbon  monoxide  there  are  3 
grams  of  carbon  combined  with  every  4  grams  of  oxygen,  and 
carbon  monoxide  is  14  times  heavier  than  hydrogen ;  its  mo- 
lecular weight  is  consequently  28.  The  atomic  weights  12  for 
carbon  and  16  for  oxygen  would  consequently  lead  us  to  write 
the  formula  for  carbon  monoxide  as  CO.  The  name  carbon 
monoxide  is  given  to  the  compound  because  it  contains  but  1 
atom  of  oxygen  in  its  molecule. 

When  carbon  is  burned  in  oxygen,  the  carbon  dioxide  formed 
is  of  exactly  the  same  volume  as  the  original  oxygen,  as  will 
appear  later.  In  other  words  :  1  volume  of  oxygen  yields  1 
volume  of  carbon  dioxide.  Accepting  Avogadro's  hypothe- 
sis, there  are  then  as  many  molecules  of  carbon  dioxide  formed 
as  there  are  molecules  of  oxygen.  The  process  may  be  ex- 
pressed by  the  equation, 

C  +  02  =  C02, 

which,  like  all  so-called  chemical  equations,  simply  expresses 
the  march  of  the  reaction.  The  sign  =  stands  for  yields* 
Some  use  the  sign  ->  instead  of  = .  The  latter  is,  however, 
more  frequently  employed. 

When  carbon  monoxide  is  burned  in  oxygen,  2  volumes  of 
the  former  unite  with  1  volume  of  the  latter  to  form  2  vol- 
umes of  carbon  dioxide.  Assuming  Avogadro's  hypothesis, 
there  must  consequently  be  formed  as  many  molecules  of 
carbon  dioxide  as  there  were  molecules  of  carbon  monoxide. 
These  relations  are  expressed  by  the  simple  equation : 

2CO  +  O2  =  2C02. 

The  above  illustrations  may  suffice  to  indicate  how  the  vapor 
densities  of  substances  have  been  employed  in  choosing  the 
atomic  weights,  the  combining  weights  having  been  ascertained 
by  careful  quantitative  analytical  or  synthetical  experiments. 


FUNDAMENTAL  LAWS   AND  THEORIES  77 

When  Avogadro  put  forth  his  hypothesis  in  1811,  it  was  by 
no  means  at  once  generally  accepted.  Indeed,  it  was  not  till 
the  vapor  densities  of  a  very  considerable  number  of  substances 
had  become  known,  that  the  value  of  the  hypothesis  was  really 
recognized.  It  was  Charles  Gerhardt,  professor  of  chemistry 
at  the  University  of  Montpellier,  who  in  1842  used  the  vapor 
densities  of  substances  as  a  guide  in  determinining  their  for- 
mulae and  in  choosing  the  atomic  weights  from  the  equivalents 
or  combining  weights,  which  were  at  that  time  in  almost 
general  use.  But  it  was  Auguste  Laurent,  professor  of  chem- 
istry at  the  University  of  Bordeaux,  who  in  1846  grasped  the 
great  value  of  Avogadro's  hypothesis  and  paved  the  way  for 
its  general  acceptance.  He  distinguished  clearly  between  atomic 
and  molecular  weights,  defining  the  molecule  as  the  smallest 
weight  of  any  substance  that  can  exist  by  itself,  and  the  atom  as 
the  smallest  weight  of  a  substance  that  can  enter  into  combination. 
But  there  are  elements  which  do  not  enter  into  compounds 
that  can  be  vaporized,  and  consequently  the  atomic  weights  of 
such  elements  cannot  be  chosen  from  the  combining  weights 
by  means  of  the  vapo-r  density,  as  described.  This  is  particu- 
larly true  of  many  of  the  metals.  In  determining  the  atomic 
weights  of  the  latter,  Berzelius  simply  took  the  number  of 
parts  by  weight  of  the  metal  that  united  with  16  parts  by 
weight  of  oxygen  as  the  atomic  weight  of  the  metal.  In  case 
a  metal  formed  more  than  one  oxide  —  recall  the  oxides  of  lead, 
for  instance  —  Berzelius  assumed  the  one  most  commonly 
found  as  containing  1  atom  of  the  metal  to  1  atom  of  oxygen, 
and  then  computed  the  formulae  of  the  other  oxides  accord- 
ingly. When  there  was  but  one  oxide  known,  as  in  the  case  of 
zinc,  for  instance,  he  assumed  that  the  molecule  consisted  of 
1  atom  of  the  metal  to  1  atom  of  oxygen.  Thus  he  proceeded 
on  the  basis  of  simplicity,  guarding  himself  by  assigning 
similar  formulae  to  substances  that  exhibit  similar  chemical 
properties.  Gerhardt,  however,  considered  it  likely  that  the 
molecules  of  the  oxides  of  the  metals  are  similar  to  the  molecule 
of  water  in  construction,  and  consequently  contain  2  atoms  of 
metal  to  1  atom  of  oxygen,  which,  of  course,  led  him  to  adopt 
atomic  weights  for  the  metals  which  were  just  half  of  the 
values  adopted  by  Berzelius.  This  led  to  considerable  discus- 
sion. But  in  1858  Stanislao  Cannizzaro,  then  professor  of 


T8 


OUTLINES   OF   CHEMISTRY 


chemistry  at  Genoa,  brought  order  into  the  confusion  that  had 
arisen  by  pointing  out  that  the  specific  heats  of  the  elements 
in  the  solid  state  may  be  employed  with  great  advantage  in 
choosing  the  true  atomic  weights  from  the  combining  weights. 
The  Law  of  Dulong  and  Petit.  —  Cannizzaro  recalled  a  simple 
relation,  discovered  by  Dulong  and  Petit  of  Paris  in  1819,  be- 
tween the  atomic  weight  of  an  element  and  its  specific  heat. 
This  relation  is  that  the  product  of  the  specific  heat  of  an  element 
in  the  solid  state  and  its  atomic  weight  is  constant.  This  law,  which 
is  known  as  the  law  of  Dulong  and  Petit,  may  also  be  expressed 
by  saying  that  the  atoms  of  the  elements  have  the  same  heat  ca- 
pacity. The  experimental  researches  of  Victor  Regnault,  the 
great  French  physicist  (1810-1878),  added  many  new  data  to 
confirm  this  law ;  but,  like  the  hypothesis  of  Avogadro,  its 
value  was  not  clearly  recognized  till  Cannizzaro  showed  how 
useful  it  is  in  fixing  the  atomic  weights  of  many  of  the  ele- 
ments. The  product  of  the  atomic  weight  and  the  specific  heat  of 
an  element  in  the  solid  state  is  approximately  6.4.  The  follow- 
ing table  gives  the  specific  heats  of  a  number  of  elements  in  the 
solid  state :  — 


ELEMENT 

SYMBOL 

ATOMIC  WEIGHT 

SPECIFIC  HEAT 

ATOMIC  HEAT 

Lithium 

Li 

6.94 

0.941 

6.53 

Sodium 

Na 

23.00 

0.293 

6.74 

Magnesium 
Aluminum 

Mg 

Al 

24.32 
27.1 

0.245 
0.214 

5.95 
5.80 

Phosphorus 
Sulphur 
Potassium 

P 

S 
K 

31.0 
32.07 
39.10 

0.202 
0.203 
0.166 

6.26 
6.50 
6.49 

Iron 

Fe 

55.85 

0.112 

6.26 

Copper 
Zinc 

Cu 

Zn 

63.57 
65.37 

0.095 
0.093 

6.04 
6.08 

Silver 
Platinum 

Ag 
Pt 

107.88 
195.2 

0.057 
0.0325 

6.15 
6.34 

Gold 

Au 

197.2 

0.0324 

6.40 

Mercury- 
Lead 

Hg 
Pb 

200.6 
207.1 

0.0333 
0.0315 

6.68 
6.52 

Glucinum 

Gl 

9.1 

0.42 

3.82 

Boron 

B 

11.0 

0.24 

2.64 

Carbon  (Graphite) 
Silicon 

C 

Si 

12.0 

28.3 

0.200 
0.177 

2.40 
5.01 

FUNDAMENTAL   LAWS   AND   THEORIES  79 

As  the  specific  heat  of  a  substance  is  a  function  of  the  tem- 
perature at  which  it  is  determined,  one  would  clearly  not  ex- 
pect the  product  of  the  atomic  weight  and  the  specific  heat  to 
yield  exactly  the  same  value.  An  inspection  of  the  table  shows 
that  the  atomic  heat  is  generally  about  6,  except  in  the  case  of 
the  last  four  elements,  where  it  varies  greatly  from  that  value. 
Now  it  is  found  that  the  specific  heats  of  glucinum,  boron,  car- 
bon, and  silicon  increase  greatly  with  rise  of  temperature,  finally 
becoming  nearly  constant.  So  the  specific  heat  of  graphite  is 
0.160  at  10°  ;  0.199  at  60° ;  0.445  at  600°  ;  and  0.460  at  900°. 
The  specific  heat  of  silicon  is  0.20  at  200°  and  0.203  at  300°; 
that  of  glucinum  is  0.617  at  400°,  and  0.620  at  500°.  Thus  at 
higher  temperatures  these  elements  approximately  obey  the  law 
of  Dulong  and  Petit,  though  at  room  temperatures  they  appar- 
ently are  exceptions. 

By  means  of  the  law  of  Dulong  and  Petit  the  atomic  weight 
of  an  element  may  be  found  by  dividing  the  atomic  heat,  ap- 
proximately 6.4,  by  the  specific  heat;  that  is, 

Atomic  Weight  = 


Specific  Heat 

This  method  can  obviously  not  be  used  for  determining  atomic 
weights  with  accuracy  ;  but  it  is  of  great  value  in  choosing  the 
true  atomic  weight  from  the  combining  weights,  and  it  is  in 
this  way  that  it  was  employed  with  great  success  by  Canniz- 
zaro.  The  latter  thus  showed  that  in  nearly  all  cases  the  values 
that  Berzelius  had  assigned  to  the  atomic  weights  of  the  metals 
were  the  correct  ones ;  and  that  only  in  a  few  instances,  like 
potassium,  sodium,  and  silver,  did  the  oxides  have  two  atoms 
of  the  metal  in  the  molecule,  though  Gerhardt  had  assumed  this 
in  all  cases.  Cannizzaro  also  showed  that  wherever  volatile 
metallic  compounds  were  known,  the  choice  of  the  atomic 
weights  from  the  vapor  density  of  these  agreed  with  the  values 
deduced  from  the  specific  heats. 

Other  Methods  of  choosing  the  Atomic  Weights  from  the  Com- 
bining Weights.  —  Many  substances  form  crystals.  Crystals 
are  solids  bounded  by  plane  faces  which  are  the  outcome  of  a 
regular  internal  structure.  Substances  which  do  riot  crystal- 
lize, that  is,  are  non-crystalline,  are  called  amorphous.  All 
crystals  may  be  classified  into  six  crystal  systems  (which  see). 


80  OUTLINES  OF   CHEMISTRY 

In  1819,  Eilhard  Mitscherlich  discovered  that  chemical  com- 
pounds which  are  similar  in  character  crystallize  in  the  same 
forms.  This  is  the  law  of  isomorphism,  for  isomorphous  sub- 
stances are  such  as  crystallize  in  the  same  form.  Whenever 
compounds  are  isomorphous,  they  are  chemically  analogous  ; 
and  so  if  the  formula  of  one  compound  has  been  determined, 
the  formulae  of  other  compounds  that  are  isomorphous  with  it 
are  readily  deduced  by  analogy.  Ismorphism  may  conse- 
quently be  used  in  choosing  the  atomic  weights  from  the  com- 
bining weights.  It  was  so  employed  by  Berzelius.  Thus,  the 
sulphate  of  magnesium  is  isomorphous  with  that  of  zinc.  If 
the  atomic  weight  of  the  latter  metal  has  been  fixed  as  65.5, 
with  the  aid  of  the  law  of  Dulong  and  Petit,  then  the  amount 
of  magnesium  that  is  required  to  replace  65.5  parts  of  zinc  by 
weight  in  the  sulphate,  namely  24.32,  is  the  atomic  weight  of 
magnesium.  In  this  way  isomorphism  has  been  of  great  use  in 
atomic  weight  determinations.  It  must  be  applied  with  great 
care,  however,  for  it  is  true  there  are  many  cases  where  com- 
pounds are  chemically  dissimilar  and  yet  possess  like  crystal 
forms.  For  this  reason,  isomorphism  is  not  so  reliable  a  guide 
as  the  specific  heat  or  vapor  density  method,  and  is  only  em- 
ployed when  other  methods-  cannot  be  used. 

In  choosing  the  atomic  weights  from  the  combining  weights, 
the  principle  of  simplicity  and  analogy  was  employed  with 
much  success  by  Berzelius.  As  already  stated,  in  the  case  of 
metals,  like  zinc  and  magnesium,  that  form  but  one  compound 
with  oxygen,  he  assumed  that  the  oxide  contains  1  atom  of 
the  metal  to  1  atom  of  oxygen.  Since  the  atomic  weight  of 
oxygen  was  taken  as  16,  the  atomic  weight  of  the  element 
combined  with  oxygen  could  readily  be  found.  Further,  Ber- 
zelius sought  to  assign  analogous  formula  to  compounds  that 
are  actually  analogous  in  their  chemical  behavior.  By  this 
method  he  consequently  chose  the  atomic  weight  of  the  ele- 
ment in  question  in  accordance  with  the  formulae  assigned. 

Finally,  the  arrangement  of  the  atomic  weights  of  the  ele- 
ments in  the  so-called  periodic  system  (which  see)  has  in  some 
cases  influenced  the  choice  of  the  atomic  weights.  .In  summary, 
then,  the  methods  of  choosing  the  atomic  weights  from  the  combin- 
ing weights  are:  (1)  the  vapor  density,  (2)  the  specific  heats  in 
the  solid  state,  (3)  isomorphism,  (4)  the  principle  of  simplicity 


FUNDAMENTAL   LAWS  AND   THEORIES 


81 


and  chemical  analogy,  and  (5)  the  periodic  system.  Concerning 
the  last  two  of  these  it  should  be  said  at  this  juncture  that  the 
manner  of  their  application  cannot  be  elucidated  before  more 
substances  Ijave  been  studied. 

Table  of  Atomic  Weights.  —  The  following  is  a  table  of  the 
atomic  weights  of  the  elements  as  adopted  by  the  International 
Commission  on  Atomic  Weights  in  1912  :  — 

INTERNATIONAL   ATOMIC   WEIGHTS,   1912 


ELEMENT 

SYMUOL 

ATOMIC 
WEIGHT 

ELEMENT 

SYMBOL 

ATOMIC 
WEIGHT 

Aluminum 

Al 

27.1 

Neodymium 

Nd 

144.3 

Antimony 

Sb 

120.2 

Neon 

Ne 

20.2 

Argon 

A 

89.88 

Nickel 

Ni 

58.68 

Arsenic 

As 

74.96 

Niton  (radium 

Barium 

Ba 

137.37 

emanation) 

Nt 

222.4 

Bismuth 

Bi 

2080 

Nitrogen 

N 

14.01 

Boron 

B 

11.0 

Osmium 

Os 

190.9 

Bromine 

Br 

79.92 

Oxygen 

0 

16.00 

Cadmium 

Cd 

112.40 

Palladium 

Pd 

106.7 

Caesium 

Cs 

132.81 

Phosphorus 

P 

31.04 

Calcium 

Ca 

40.07 

Platinum 

Pt 

195.2 

Carbon 

C 

12.00 

Potassium 

K 

39.10 

Cerium 

Ce 

140.25 

Praseodymium 

Pr 

140.6 

Chlorine 

Cl 

35.46 

Radium 

Ra 

226.4 

Chromium 

Cr 

52.0 

Rhodium 

Rh 

102.9 

Cobalt 

Co 

58.97 

Rubidium 

Rb 

85.45 

Columbium 

Cb 

93.5 

Ruthenium 

Ru 

101.7 

Copper 

Cu 

63.57 

Samarium 

Sa 

150.4 

Dysprosium 

r>y 

162.5 

Scandium 

Sc 

44.1 

Erbium 

Er 

167.7 

Selenium 

Se 

79.2 

Europium 

Eu 

152.0 

Silicon 

Si 

28.3 

Fluorine 

F 

19.0 

Silver 

Ag 

107.88 

Gadolinium 

Gd 

157.3 

Sodium 

N! 

23.00 

Gallium 

Ga 

69.9 

Strontium 

Sr 

87.63 

Germanium 

Ge 

72.5 

Sulphur 

S 

32.07 

Glucinum 

Gl 

9.1 

Tantalum 

Ta 

181.5 

Gold 

Au 

197.2 

Tellurium 

Te 

127.5 

Helium 

He 

3.99 

Terbium 

Tb 

159.2 

Hydrogen 

H 

1.008 

Thallium 

Tl 

204.0 

Indium 

In 

114.8 

Thorium 

Th 

232.4 

Iodine 

I 

126.92 

Thulium 

Tm 

168.5 

Iridium 

Ir 

193.1 

Tin 

Sn 

119.0 

Iron 

Fe 

55.84 

Titanium 

Ti 

48.1  " 

Krypton 

Kr 

82.92 

Tungsten 

W 

184.0 

Lanthanum 

La 

139.0 

Uranium 

U 

238.5 

Lead 

Pb 

207.10 

Vanadium 

V 

51.0 

Lithium 

Li 

6.94 

Xenon 

Xe 

130.2 

Lutecium 

Lu 

174.0 

Ytterbium 

Magnesium 
Manganese 

Mg 
Mn 

24.32 
54.93 

(Neoytterbium) 
Yttrium 

Yb 
Yt 

172.0 

89.0 

Mercury 

Hg 

200.6 

Zinc 

Zn 

65.37 

Molybdenum 

Mo 

96.0 

Zirconium 

Zr 

90.6 

82  OUTLINES  OF  CHEMISTRY 

Interpretation  of  a  Chemical  Formula.  —  A  chemical  formula 
expresses  (1)  what  elements  occur  in  the  compound,  (2)  the 
relative  weights  in  which  these  elements  occur  in  the  compound, 
and  (3)  the  weight  of  22.38  liters  of  the  vapor  ^of  the  com- 
pound under  standard  conditions.  In  case  the  compound  can- 
not be  converted  into  the  vapor  state,  the  formula  is  derived 
from  a  study  of  the  freezing  or  boiling  point  of  its  solution, 
the  crystalline  form,  the  specific  heat,  or  from  its  chemical 
behavior  and  analogy  to  other  compounds.  These  facts  are 
thus  recorded  in  the  formula.  So  the  formula  of  carbonic 
acid  gas  CO2  tells  us  that  this  compound  consists  of  carbon 
and  oxygen  in  the  proportions  of  12  parts  of  the  former  to  32 
parts  of  the  latter  by  weight.  It  also  tells  that  •  the  weight  of 
22.38  liters  of  the  gas  under  standard  conditions  is  44  grams, 
or  that  the  gas  is  22  times  as  heavy  as  hydrogen.  Thus  we 
see  that  chemical  formulae  are  a  system  of  shorthand  writing, 
as  it  were,  for  they  express  in  a  small  space  the  salient  facts 
known  about  a  compound. 

Valence  and  Structural  Formulae.  —  For  hydrochloric  acid  we 
have  developed  the  formula  HC1.  In  this  compound  1  atom 
of  hydrogen  is  combined  with  1  of  chlorine.  In  water  H2O, 
on  the  other  hand,  2  atoms  of  hydrogen  are  combined  with  1 
of  oxygen.  The  power  which  an  atom  of  one  element  has  to 
unite  with  one  or  more  atoms  of  other  elements  is  called  its 
valence.  Thus  in  hydrochloric  acid,  hydrogen  and  chlorine 
each  have  a  valence  of  one.  Hydrochloric  acid  is  said  to  be  a 
saturated  compound,  for  it  will  unite  with  neither  more  hydro- 
gen nor  more  chlorine.  Hydrogen  always  has  a  valence  of 
one,  it  is  consequently  called  a  univalent  element  or  a  monad. 
The  number  of  hydrogen  atoms,  or  other  univalent  atoms,  with 
which  an  atom  of  a  given  element  combines  determines  the  valence 
of  the  latter.  In  water,  we  have  2  hydrogen  atoms  united  with 
1  oxygen  atom.  Oxygen,  consequently,  has  a  valence  of  2 ; 
i.e.  it  is  a  bivalent  element  or  dyad. 

The  formula  for  water  is  also  written  H  —  O  —  H  to 
indicate  that  each  of  the  atoms  of  hydrogen  is  bound  to 
oxygen,  which  idea  may  be  deduced  from  the  fact  that 
when  sodium  acts  on  water  only  half  of  the  hydrogen 
of  the  latter  is  displaced,  and  that  the  other  half  of  the  hy- 
drogen is  set  free  when  the  resulting  sodium  hydroxide 


FUNDAMENTAL  LAWS   AND   THEORIES  83 

is  treated  with  zinc,  as  indicated  by  the  equations  (compare 
Chapter  II)  — 

H-O-H+     Na     =  NaOH  +        H. 

Water      +  sodium  =  sodium  hydroxide  -h  hydrogen. 

2  NaOH        +     Zn     =         Na2O2Zn          +        H2. 
Sodium  hydroxide  +    zinc     =     sodium  zincate  +  hydrogen. 

The  formula  H  —  O  —  H  is  therefore  the  structural  formula  for 
water.  It  is  clear  that  it  is  derived  from  reactions  that  water 
will  undergo,  and  is  consequently  merely  a  brief  way  of  express- 
ing these  changes.  Structural  formulae  are  often  used  in  chem- 
istry, particularly  in  connection  with  the  compounds  of  carbon. 
It  is  not  to  be  thought  that  such  a  formula  expresses  the 
actual  conditions  that  exist  within  the  molecule  itself;  it  is 
rather  simply  a  concise  expression  of  the  reactions  which  the 
compound  in  question  will  undergo  with  other  chemical 
compounds. 

In  chlorine  monoxide  C12O  chlorine  is  univalent  and  oxy- 
gen is  bivalent.  The  structural  formula  of  the  compound  is 
Cl  —  O  —  Cl ;  we  may  consider  it  as  water  in  which  the  hydro- 
gen atoms  are  replaced  by  chlorine.  In  chlorine  dioxide  C1O2 
the  oxygen  is  bivalent,  and  the  chlorine  has  a  valence  of  four ; 
i.e.  it  is  quadrivalent,  the  formula  being  O  =  Cl  =  O.  Oxygen 
is  always  bivalent  except  in  very  rare  cases.  It  is  thus  clear 
that  the  number  of  oxygen  atoms,  or  other  atoms  of  known  valence, 
with  which  an  atom  of  another  element  combines,  may  also  serve  to 
ascertain  the  valence  of  the  latter.  In  carbon  dioxide  CO2  carbon 
is  quadrivalent ;  thus,  O  =  C  =  O.  In  carbon  monoxide  CO 
carbon  is  bivalent;  thus,  C  =  O  ;  or  sometimes  it  is  considered 
as  quadrivalent,  two  combining  powers  or  bonds  being  free  or 
unsaturated,  thus,  =  C  =  O.  This  at  once  brings  us  to  the 
question  whether  the  valence  of  an  element  is  always  the  same 
or  not.  There  has  been  considerable  dispute  over  this  question, 
but  now  it  is  quite  generally  held  that  the  valence  of  an  ele- 
ment may  vary  in  different  compounds.  The  highest  valence 
which  an  element  exhibits  in  any  known  compound  is  called 
its  maximum  valence. 

The  valence  of  an  element  may  vary  from  one  to  eight, 
though  in  case  of  most  elements  it  varies  but  slightly.  As 
already  stated,  hydrogen  is  always  univalent ;  oxygen  is  almost 


84  OUTLINES   OF  CHEMISTRY 

always  bivalent;  and  carbon  may  practically  always  be  con- 
sidered as  quadrivalent,  though  in  some  compounds  it  is  biva- 
lent and  even  trivalent.  Chlorine  is  always  univalent  toward 
hydrogen,  while  toward  oxygen  it  may  be  either  univalent, 
bivalent,  quadrivalent,  or  heptavalent.  The  valences  of  the 
various  elements  will  be  taken  up  in  connection  with  the 
description  of  each  element,  for  the  subject  cannot  be  con- 
sidered fully  except  in  connection  with  actual  illustrations. 

Nomenclature. — The  names  of  the  metallic  elements  end  in 
urn,  like  lithium,  sodium,  bariwm,  etc.,  except  in  case  of  some 
metals  that  have  been  known  for  a  very  long  time,  which 
retain  their  old  names,  as  iron,  lead,  gold,  silver,  etc.  The 
elements  selenium  and  tellurium  are  not  metals.  They  were 
thought  to  be  such  when  discovered,  on  account  of  their 
outward  properties,  hence  the  ending  um  in  their  names. 

Substances  containing  but  two  elements  are  called  binary 
compounds ;  their  names  end  in  ide.  Thus,  common  salt  NaCl 
is  sodium  chloride  ;  magnesia  is  magnesium  oxide  MgO  ;  lime 
is  calcium  oxide  CaO,  etc.  In  some  cases  the  suffix  ide  is  used 
in  connection  with  compounds  containing  more  than  two  ele- 
ments. In  all  these,  however,  two  or  more  of  the  elements  act 
as  a  group,  i.e.  a  unit  or  radical,  so  called,  which  may  pass 
from  compound  to  compound  ;  thus,  sodium  hydroxide  NaOH 
contains  the  OH  group,  which  is  called  the  Tiydroxyl  group. 
This  OH  group  passes  from  one  compound  to  another  as  a  unit, 
and  when  elements  or  other  groups  are  combined  with  this 
group,  the  compounds  formed  are  termed  hydroxides.  So  we 
have  calcium  hydroxide  Ca(OH)2  formed  when  calcium  acts 
on  water, 

Ca  +  2H20  =  Ca(OH)2+H2; 

or  when  lime,  calcium  oxide  CaO,  is  treated  with  water,  i.e.  is 
slaked,  thus  :  — 

CaO  +  H2O  =  Ca(OH)2. 

In  the  hydroxyl  group  OH  one  of  the  two  combining  powers 
or  bonds  of  oxygen  is  satisfied  by  hydrogen,  the  other  bond 
being  free.  The  group  is  therefore  univalent,  and  we  may 
write  it  thus  :  —  O  —  H. 

When  two  elements  form  more  than  one  compound  with  each 
other  (which  is  frequently  the  case),  a  name  indicating  the  num- 


FUNDAMENTAL  LAWS  AND  THEORIES  85 

ber  of  atoms  of  the  one  element  that  are  united  with  the  other 
is  given  to  each  compound.  Thus  we  have  carbon  monoxide 
CO,  carbon  dioxide  CO2,  sulphur  dioxide  SO2,  sulphur  trioxide 
SO3,  phosphorus  trichloride  PC13,  phosphorus  pentachloride 
PC16,  lead  sesquioxide  Pb2O3.  The  ending  ous  is  frequently 
used  for  one  compound  and  the  ending  ic  for  another  compound 
richer  than  the  former  in  one  of  the  ingredients.  Thus  SO2  or 
sulphur  dioxide  is  also  called  sulphurous  oxide,  and  SO3  or  sul- 
phur trioxide  is  also  called  sulphuric  oxide.  Again,  PC13  is 
phosphorus  trichloride  or  phosphorous  chloride,  and  PC16  is 
phosphorus  pentachloride  or  phosphoric  chloride.  When  more 
than  two  compounds  are  formed  by  two  elements,  the  endings 
ous  and  ic  are  retained  and  the  prefixes  proto,  hypo  or  sub,  and 
per  are  added  as  required.  Thus  litharge  PbO  is  lead  mon- 
oxide,  lead  protoxide  or  plumbic  oxide ;  black  oxide  of  lead 
Pb2O  is  lead  swfoxide  or  plumbous  oxide  ;  Pb2O3is  lead  sesqui- 
Oxide  ;  minium  or  red  lead  Pb3O4  is  the  proto-sesquioxide  of 
lead,  i.e.  PbO  •  Pb2O3;  brown  oxide  of  lead  PbO2  is  lead  dioxide 
or  lead  peroxide.  The  prefix  per  stands  for  the  highest  oxida- 
tion stage  in  the  case  of  oxides,  for  the  highest  chlorination 
stage  in  the  case  of  chlorides,  etc.  Chlorine  monoxide  or  pro- 
toxide C12O  is  also  called  %^ochlorous  oxide.  Water  H2O  is 
hydrogen  protoxide  or  monoxide,  or  hydrogen  hydroxide,  or 
hydroxyl  hydride.  The  prefix  hypo  is  rarely  used  in  case  of 
binary  compounds. 

In  ternary  compounds,  that  is,  those  that  are  made  up  of  three 
elements,  somewhat  similar  designations  are  employed,  which 
will  be  explained  when  compounds  of  this  character  are  con- 
sidered. 

Chemical  Equations.  Retrospect.  —  Chemical  compounds  may 
be  designated  by  means  of  symbols  or  formulae,  as  we  have 
seen,  and  chemical  changes  may  be  indicated  by  writing  equa- 
tions in  which  these  formulae  are  used  instead  of  the  names  of 
the  compounds.  Reviewing  the  work  on  hydrogen,  oxygen, 
and  chlorine,  and  writing  the  principal  chemical  changes  that 
have  been  studied  in  form  of  chemical  equations,  we  have  as 
follows :  — 

(1)  Preparation  of  Hydrogen 

Na  +  H2O  =  NaOH  +        H. 

Sodium        -f  water  =  sodium  hydroxide  •+•  hydrogen 


86  OUTLINES  OF  CHEMISTRY 

H2SO4  +    Zn    =         ZnSO4  +       H2. 

Sulphuric  acid      -f  zinc   =     zinc  sulphate      +  hydrogen. 

2KOH  +    Zn    =          K202Zn          +       H2. 

Potassium  hydroxide  +  zinc    =  potassium  zincate  4-  hydrogen. 

3KOH  +    Al    =         K303A1  +       3  H. 

Potassium  hydroxide  -f-  aluminum 

=  potassium  aluminate-h  hydrogen. 

2  H2O  (on  electrolysis)  =          2  H2  +      O2. 

3Fe  +4H2O     =          Fe3O4  +     4  H2. 

Mg  +H20       =         MgO  +      H2. 

(2)  Preparation  of  Oxygen 

HgO        (on  heating)   =        Hg  4.     O. 

Mercuric  oxide  (on  heating)  =    mercury  -f  oxygen. 

Ag2O       (on  heating)   =      2  Ag  +     O. 

Argentic  oxide  (on  heating)  =      silver  +  oxygen. 

KC1O3      (on  heating)  =      KC1  +     3  O. 

Potassium  chlorate  (on  heating) 

=  potassium  chloride  +  oxygen. 

3  MnO2  (on  ignition)       =      Mn3O4  +      O2. 

Manganese  dioxide  (on  ignition) 

=  manganese  proto-sesquioxide  +  oxygen. 

KNO3  (on  heating)          =      KNO2  +      O. 

Potassium  nitrate  or  saltpeter  (on  heating) 

=  potassium  nitrite  +  oxygen. 

(3)  Oxidations 

2H2+     02=2H20. 
Mg  +    O  =  MgO. 
Cu  +    O  =  CuO. 


phosphorus  pentoxide. 
C  +  02    =C02. 
3  Fe  +  2  O2  =  Fe3O4. 


2Fe  +  3O  =Fe2O3, 


ferric  oxide  or  sesquioxide  of  iron 


S  +  O2  =  SOa. 


FUNDAMENTAL  LAWS  AND   THEORIES  87 

(4)  Reductions 

CuO    +     H2  =  Cu     +  H20. 
Fe8O4  +  4  H2  =  3  Fe  +  4  H2O. 


(5)  Preparation  of  Chlorine 

MnCl2  + 

manganous  chloride. 


4HCl  +  MnO2   =  MnCl2 


(6)  Reactions  of  Chlorine 

H20+     C12=2HC1  +  O. 
P  +  3  Cl  =  PC18, 

phosphorus  trichloride. 
P  +  5  Cl  =  PC16, 

phosphorus  pentachloride. 
Sb  +  3  Cl  =  SbCl3, 

antimony  trichloride. 
CioHi8     +  8  C12  =  16  HC1  +  10  C. 
turpentine. 

Phenomena  of  the  Nascent  State.  —  When  a  dilute  solution  of 
sulphuric  acid  is  acting  on  zinc,  the  hydrogen  liberated  will 
reduce  many  substances  like  potassium  permanganate,  potas- 
sium bichromate,  or  saltpeter,  if  these  are  added  directly  to  the 
mixture  in  the  generator.  The  reduction  will  not  take  place 
if  the  hydrogen  is  passed  through  solutions  of  these  salts 
contained  in  a  separate  vessel.  The  explanation  of  this  as 
commonly  given  is  that  at  the  moment  of  liberation,'  the 
hydrogen  is  in  the  so-called  nascent  state,  i.e.  in  an  atomic  con- 
dition represented  by  H,  whereas  afterwards  it  passes  over  into 
the  molecular  condition  H2.  While  in  the  nascent  state  the 
hydrogen  is  more  active  than  in  the  molecular  state,  and  it  con- 
sequently effects  many  reductions.  Similarly  we  may  have 
nascent  oxygen  O,  as  compared  with  molecular  oxygen  O2. 
Cases  of  this  kind  will  be  mentioned  later. 


CHAPTER  VII 

OZONE,   ALLOTROPY,   AND   HYDROGEN   PEROXIDE 

History,  Occurrence,  and  Preparation  of  Ozone.  —  When  a  fric- 
tional  electrical  machine  is  operated,  there  is  observed  in  its 
neighborhood  a  peculiar  characteristic  odor,  which  is  sometimes 
described  as  similar  to  the  odor  of  chlorine,  burnt  sulphur,  or 
garlic.  The  observation  that  this  smell  is  produced  when 
electric  sparks  are  passed  through  oxygen  was  made  in  1785 
by  Van  Marum,  who  had  constructed  an  especially  powerful 
machine.  The  same  odor  is  noticed  whenever  electric  sparks 
pass  through  the  air,  as,  for  instance,  from  an  induction  coil,  or 
when  objects  are  struck  by  lightning.  In  1840  Christian 
Schonbein,  professor  at  the  University  of  Basel,  showed  that 
when  water  is  electrolyzed  the  oxygen  obtained  always  con- 
tains some  of  this  odoriferous  substance,  which  he  named  ozone, 
meaning  a  smell.  From  the  fact  that  ozone  is  produced  when 
electric  sparks  pass  through  pure,  dry  oxygen,  it  is  clear  that 
the  substance  consists  of  oxygen.  By  means  of  the  silent 
electrical  discharge  ozone  is  produced  in  larger  quantities. 
For  this  purpose  an  apparatus  like  that  in  Fig.  28  is  commonly 
employed.  The  apparatus  is  blown  of  one  piece  of  glass.  The 


OZONE,   ALLOTROPY,   AND   HYDROGEN  PEROXIDE  89 

outside  of  the  tube  A  is  coated  with  tin  foil,  as  is  also  the 
inside  of  the  tube  B,  as  indicated.  Dry  oxygen  is  passed 
through  the  apparatus  as  shown,  and  when  the  tin  foil  coatings 
are  connected  with  the  poles  of  an  induction  coil,  ozone  issues 
at  0.  In  this  way  about  5  to  8  per  cent  of  the  oxygen  is  con- 
verted into  ozone.  By  liquefying  oxygen  and  ozone  by  means 
of  liquid  air,  a  liquid  is  obtained  which  upon  slow  evaporation 
leaves  a  very  dark  blue  liquid  consisting  of  about  86  per  cent 
ozone  and  14  per  cent  oxygen. 

Besides  being  formed  by  means  of  electrical  discharges  and 
in  the  electrolysis  of  water,  ozone  is  produced  in  chemical 
reactions,  notably  when  moist  phosphorus  slowly  oxidizes  in 
the  air ;  also  generally  when  oxygen  is  rapidly  evolved,  as  by 
heating  potassium  chlorate,  or  when  potassium  permanganate 
is  treated  with  strong  sulphuric  acid.  Further,  ozone  is  formed 
in  very  small  quantities  when  hydrogen  burns  in  oxygen.  By 
the  action  of  fluorine  on  water,  oxygen  containing  up  to  15 
per  cent  of  ozone  is  formed. 

Relation  between  Ozone  and  Oxygen,  Allotropy.  —  As  already 
stated,  ozone  is  produced  from  oxygen.  By  passing  ozone 
through  a  red-hot  tube,  it  is  again  converted  into  oxygen. 
Under  standard  conditions,  22.38  liters  of  ozone  weigh  48 
grams.  The  molecular  weight  of  ozone  is  consequently  48; 
and  since  the  atomic  weight  of  oxygen  is  16,  the  formula  of 
ozone  is  O3.  The  change  of  oxygen  to  ozone  is  expressed  by 
the  following  equation :  — 

3  02  (plus  energy)  =  2  O3. 

The  energy  that  must  be  added  to  oxygen  to  convert  it  into 
ozone  may  be  obtained  from  the  silent  electric  discharge,  or 
from  chemical  changes,  as  we  have  seen.  When  ozone  is 
heated,  the  reaction  is  reversed.  We  have  here  then  a  reversi- 
ble reaction.  This  fact  may  be  expressed  thus :  — 

3  O2  (plus  energy)  5*  2  O3, 

where  the  arrows  are  used  instead  of  the  usual  sign  of  equality. 
In  forming  ozone,  oxygen  shrinks  from  3  volumes  to  2  vol- 
umes and  simultaneously  a  considerable  amount  of  energy  is 
absorbed.  Ozone  is  called  an  allotropic  form  of  oxygen.  The 
property  which  some  elements  possess  of  occurring  in  two  or 
more  forms  is  called  allotropy. 


90  OUTLINES   OF   CHEMISTRY 

Ozone  is  a  much  more  powerful  oxidizing  agent  than 
oxygen.  Many  of  the  reactions  which  take  place  in  oxygen 
only  at  higher  temperatures  proceed  readily  in  ozone  at  room 
temperatures. 

Properties  of  Ozone.  —  In  thick  layers  ozone  gas  has  a  bluish 
color.  Inhaled  in  quantity,  it  attacks  the  mucous  membranes 
and  produces  headache.  Liquid  ozone  is  indigo-blue  in  color, 
and  boils  at  —  119°  under  atmospheric  pressure.  The  liquid  is 
strongly  magnetic.  On  warming,  it  is  liable  to  explode,  due 
to  sudden  change  of  the  substance  to  ordinary  oxygen.  Ac- 
cording to  Ladenburg,  1000  volumes  of  water  dissolve  10  vol- 
umes of  ozone.  It  acts  slowly  on  water,  forming  oxygen  and 
hydrogen  peroxide  (which  see),  and  the  solubility  in  water  may 
be  due  to  this  fact. 

The  chief  chemical  property  of  ozone  is  its  oxidizing  poiver.  It 
will  bleach  litmus,  indigo,  and  other  dyestuffs,  the  colors  being 
destroyed  by  oxidation.  Ozone  destroys  disease  germs  and 
other  minute  organisms,  and,  consequently,  it  is  used  as  a 
germicide  in  sterilizing  drinking  water.  Ozone  is  soluble  in 
turpentine,  also  in  oil  of  cinnamon  and  other  similar  oils.  In 
solutions,  ozone  is  also  a  powerful  oxidizing  agent.  On  account 
of  its  oxidizing  power,  it  causes  many  oils  to  thicken  and  become 
resinous. 

Ozone  rapidly  oxidizes  such  substances  as  silver,  lead,  arsenic, 
phosphorus,  and  sulphur,  to  their  highest  stages  of  oxidation. 
It  is  the  most  powerful  oxidizing  agent  known.  It  acts  on 
potassium  iodide  solutions,  liberating  iodine,  thus :  — 

2  KI  +  H2O  +  O3  =  2  KOH  +  O2  +  I2. 

Iodine  turns  starch  paste  blue,  and  so  when  a  strip  of  paper 
saturated  with  starch  paste  plus  a  solution  of  potassium  iodide 
is  exposed  to  ozone,  the  paper  turns  deep  blue  in  color.  This 
is  a  common  test  for  ozone.  However,  it  must  be  used  with 
proper  care ;  for,  as  we  shall  see,  there  are  other  things  besides 
ozone  that  turn  starch  potassium  iodide  paper  blue.  The  above 
reaction  may  be  used  in  estimating  the  amount  of  ozone  in  a 
given  sample  of  oxygen,  by  determining  the  quantity  of  iodine 
set  free.  Ozone  is  the  only  gaseous  oxidizing  agent  that  will 
blacken  a  bright  silver  foil,  and  consequently  this  test  is  used 
in  detecting  ozone  in  presence  of  other  oxidizing  gases. 


OZONE,    ALLOTROPY,    AND    HYDROGEN   PEROXIDE  91 

History,  Occurrence,  and  Preparation  of  Hydrogen  Peroxide.  - 

In  1818  Thenard  prepared  a  compound  of  hydrogen  and  oxy- 
gen containing  twice  as  much  oxygen  as  there  is  in  water.  He 
treated  barium  dioxide  with  hydrochloric  acid,  thus  :  — 

Ba02  +  2  HC1  =  BaCl2  +  H2O2. 

Both  the  barium  chloride  and  hydrogen  peroxide  (which  is 
also  called  hydrogen  dioxide  or  hydroperoxide)  remain  in  solu- 
tion. Hydrogen  peroxide  may  also  be  prepared  by  adding 
barium  dioxide  to  cold,  dilute  sulphuric  acid  :  — 

Ba02  +  H2S04  =  BaS04  +  H2O2  ; 

or  by  passing  carbon  dioxide  through  water  and  gradually 
adding  barium  dioxide  in  small  amounts :  — 

BaO2  +  CO2  +  H2O  =  BaCO3  +  H2O2. 

Barium  sulphate  and  barium  carbonate  are  insoluble  in  water, 
and  hence  may  be  removed  by  nitration ;  and  thus  a  nitrate, 
which  is  an  aqueous  solution  of  hydrogen  peroxide,  may  be 
obtained. 

When  ozone  acts  on  water,  hydrogen  peroxide  is  produced:  — 

H2O  +  O3  =  H2O2  +  O2. 

Hydrogen  peroxide  occurs  in  very  small  amounts  in  the  air, 
and  this  is  probably  due  to  the  fact  that  ozone  has  been  pro- 
duced, which  in  turn  has  acted  on  the  moisture  in  the  air.  It 
is  consequently  very  doubtful  whether  ozone  itself  occurs  in 
air.  It  should  be  stated  here  that  the  occurrence  of  hydrogen 
peroxide  in  the  air  has  been  questioned  by  some  chemists,  the 
claim  being  made  that  the  strong  oxidations  observed  may  very 
well  be  caused  by  oxides  of  nitrogen  which  are  present  in  the 
atmosphere. 

Hydrogen  peroxide  may  also  be  formed  by  treating  cold, 
dilute  hydrochloric  acid  with  sodium  peroxide :  — 

2  HC1  +  Na202  =  2  NaCl  +  H2O2. 

Both  the  sodium  chloride  and  hydrogen  peroxide  remain  in 
solution.  Insj^ad  of  the  peroxide  of  barium  or  sodium,  that 
of  potassium  or  strontium  may  be  used.  By  distilling  an 
aqueous  solution  of  hydrogen  peroxide  in  a  partial  vacuum,  the 
water  passes  off  first,  leaving  hydrogen  peroxide  in  the  retort. 


92  OUTLINES   OF  CHEMISTRY 

On  heating  a  3  per  cent  solution  of  hydrogen  peroxide  on  the 
water  bath  to  temperatures  below  70°,  in  a  retort  from  which 
the  air  has  been  exhausted  so  as  to  create  a  partial  vacuum,  a 
45  per  cent  solution  may  readily  be  obtained  without  loss.  On 
continuing  the  distillation  further,  nearly  pure  hydrogen  perox 
ide  passes  over  between  84°  and  85°  at  68  mm.  pressure. 

Properties  of  Hydrogen  Peroxide.  — Pure  hydrogen  peroxide 
is  a  colorless,  sirupy  liquid,  which,  like  water,  has  a  bluish 
hue  in  thick  layers.  At  0°  its  specific  gravity  is  1.458.  It 
boils  at  69°  under  26  mm.  pressure,  and  at  84°  to  85°  under 
68  mm.  pressure.  It  forms  colorless  prismatic  crystals  which 
melt  at  -  2°. 

Hydrogen  peroxide  slowly  decomposes  into  water  and  oxygen 
on  standing.  In  the  sunlight  the  decomposition  proceeds  more 
rapidly.  By  warming  hydrogen  peroxide  the  rate  of  decom- 
position is  increased  ;  and  at  100°  the  evolution  of  oxygen 
becomes  so  rapid  as  to  cause  explosion.  It  is,  therefore,  neces- 
sary to  distill  hydrogen  peroxide  in  a  vacuum,  so  that  it  will 
not  need  to  be  heated  to  a  temperature  at  which  violent  decom- 
position sets  in. 

Solutions  of  hydrogen  peroxide  have  a  peculiar  bitter,  disa- 
greeable taste.  Concentrated  solutions  act  on  the  skin.  The 
aqueous  solutions  on  the  market  usually  contain  about  3  per 
cent  hydrogen  peroxide,  though  30  per  cent  solutions  are  also 
now  placed  on  sale.  The  latter  are  kept  in  small  bottles  coated 
with  paraffin  on  the  inside;  for  in  contact  with  glass  the  solu- 
tion soon  suffers  decomposition  on  account  of  the  fact  that 
alkali  is  dissolved  from  the  glass. 

In  contact  with  platinum  black,  manganese  dioxide,  or  finely 
divided  silver,  gold,  or  carbon,  hydrogen  peroxide  is  decomposed 
into  oxygen  and  water  even  at  room  temperatures  and  in  dilute 
solutions.  The  action  is  more  rapid  at  higher  temperatures. 
All  these  cases  are  illustrations  of  catalytic  or  contact  action. 

,  Hydrogen  peroxide  is  2,  strong  oxidizing  agent.  It  will  act 
on  black  sulphide  of  lead  and  convert  it  into  lead  sulphate, 
which  is  a  white  salt :  — 

PbS  +  4  H2O2  =  PbS04  +  4  H20. 
Potassium  iodide  in  solution  is  oxidized  thus :  — 
2KI  +  H202=2KOH-i-I2. 


OZONE,   ALLOTROPY,   AND   HYDROGEN  PEROXIDE  93 

For  this  reason,  starch  potassium  iodide  paper  may  be  used  to 
detect  the  presence  of  hydrogen  peroxide.  The  action  goes  on 
much  more  slowly  than  in  the  case  of  ozone ;  but  the  addition 
of  a  little  ferrous  sulphate  hastens  the  action  very  markedly,  so 
that  the  test  is  really  a  sensitive  one.  In  presence  of  ozone, 
which  also  liberates  iodine  from  potassium  iodide,  this  test  for 
hydrogen  peroxide  can,  of  course,  not  be  used.  Hydrogen 
peroxide  does  not  oxidize  a  bright  silver  foil  as  ozone  does,  and 
thus  the  latter  may  be  detected  in  presence  of  the  former. 

In  contact  with  blood,  meat,  and  the  mucous  membranes, 
hydrogen  peroxide  decomposes.  The  oxygen  thus  liberated 
destroys  germs  by  oxidizing  them,  hence  the  use  of  hydrogen 
peroxide  in  medicine  as  a  gargle  and  an  antiseptic. 

When  limewater  is  treated  with  hydrogen  peroxide  solution, 
a  precipitate  of  calcium  peroxide  is  formed :  — 

Ca(OH)2  +  H2O2  =  CaO2  +  2  H2O. 

The  action  on  the  hydroxide  of  barium  or  strontium  is  similar. 
All  of  these  peroxides  may  be  regarded  as  hydrogen  peroxide 
in  which  the  hydrogen  is  replaced  by  metals. 

When  hydrogen  peroxide  solution  is  slightly  acidified  with 
sulphuric  acid,  and  a  few  drops  of  potassium  bichromate  solu- 
tion and  some  ether  are  added,  and  the  mixture  is  then  shaken, 
an  indigo -blue  compound  is  formed  which  dissolves  in  the  ether 
and  so  finally  collects  in  the  light  ethereal  layer  on  standing. 
This  reaction  is  used  as  a  test  for  hydrogen  peroxide.  The 
nature  of  the  blue  compound  is  not  known  with  certainty, 
though  it  is  probably  perchromic  acid. 

While  hydrogen  peroxide  is  an  oxidizing  agent,  it  may  also  at 
times  act  as  a  reducing  agent,  in  which  case  ordinary  oxygen  gas 
is  evolved.  So  the  oxides  of  metals  like  silver,  gold,  and  plati- 
num suffer  reduction  to  the  metallic  state  when  treated  with 
hydrogen  peroxide,  thus  :  — 

Ag20  +  H202  =  2  Ag  +  H20  +  02. 

We  see  that  hydrogen  peroxide  in  such  cases  loses  one  atom  of 
oxygen  which  unites  with  oxygen  of  the  metallic  oxides  and 
escapes  as  ordinary  oxygen  gas.  Lead  peroxide  is  changed  to 
lead  monoxide :  — 

PbO2  +  H2O2  =  PbO  +  H2O  +  O2. 
Added  to  a  potassium  permanganate  solution  acidified  with  sul* 


94  OUTLINES   OF  CHEMISTRY 

phuric  acid,  hydrogen  peroxide  reduces  the  permanganate  with 
liberation  of  oxygen  and  formation  of  a  solution  of  potassium 
sulphate  and  manganous  sulphate,  which  is  nearly  colorless  :  — 


This  reaction  is  used  in  the  quantitative  determination  of  the 
strength  of  hydrogen  peroxide  solutions  ;  for  if  a  certain  volume 
of  a  potassium  permanganate  solution  of  known  strength  is  just 
decolorized  by  a  known  volume  of  a  hydrogen  peroxide  solu- 
tion, the  strength  of  the  latter  can  readily  be  computed  from 
the  data  given  in  the  above  equation. 

It  would  seem  rather  peculiar  that  hydrogen  peroxide,  which 
is  a  good  oxidizing  agent,  may  also  serve  in  effecting  reductions. 
It  must  be  borne  in  mind,  however,  that  it  only  reduces  com- 
pounds that  are  rich  in  oxygen  which  is  readily  set  free.  The 
explanation  of  the  reduction  is  that  when  compounds  like  potas- 
sium permanganate,  or  oxides  of  silver,  gold,  lead,  etc.,  are 
brought  in  contact  with  hydrogen  peroxide,  the  tendency  to 
form  the  ordinary  oxygen  molecule  O2,  that  is,  the  attraction  of 
oxygen  for  oxygen,  is  so  great  that  the  compounds  mutually 
reduce  each  other. 

Formula  of  Hydrogen  Peroxide.  —  Thenard,  the  discoverer  of 
hydrogen  peroxide,  determined  that  it  consists  of  16  parts  of 
oxygen  to  1  part  of  hydrogen  by  weight.  The  simplest  for- 
mula one  could  assign  to  the  compound  would  therefore  be 
HO,  the  atomic  weight  of  oxygen  being  16.  However,  the 
fact  that  water  H2O  and  oxygen  are  formed  when  hydrogen 
peroxide  decomposes,  is  much  better  indicated  by  adopting  the 
formula  H2O2  for  the  latter  substance.  The  vapor  density  of 
hydrogen  peroxide  cannot  well  be  determined  because  the  sub- 
stance is  so  unstable,  and  so  the  weight  of  22.38  liters  of  its 
vapor  under  standard  conditions  is  unknown.  Its  molecular 
weight  has,  however,  been  found  to  be  34,  from  a  study  of  the 
freezing  point  of  its  aqueous  solution. 

The  fact  that  hydrogen  peroxide  decomposes  into  water  and 

TT\ 

oxygen  has  led  Kingsett  to  ascribe  to  it  the  formula  T    /O  =  O, 

in  which  it  will  be  seen  that  one  oxygen  atom  is  regarded  as  a 
tetrad  and  the  other  as  a  dyad.  From  a  study  of  the  index  of 
refraction  of  the  substance,  Briihl  has  on  the  other  hand  sug- 
gested that  both  oxygen  atoms  are  tetrads  and  that  the  formula 


OZONE,   ALLOTROPY,   AND   HYDROGEN  PEROXIDE  95 

should  be  written,  H  —  O  =  O  —  H.  As  a  rule  chemists  regard 
both  atoms  of  oxygen  as  bivalent,  writing  the  structural  for- 
mula of  hydrogen  peroxide,  H  —  O  —  O  —  H. 

A  structural  formula  expresses  not  only  the  qualitative  and 
quantitative  composition  of  a  substance  and  its  molecular  weight, 
but  it  also  indicates  its  chemical  behavior.  This  is  accomplished 
by  arranging  the  relative  position  of  the  atoms  in  the  formula  so 
as  to  indicate  what  chemical  changes  the  compound  will  undergo. 

Uses  of  Hydrogen  Peroxide.  —  As  already  stated,  hydrogen 
peroxide  is  used  in  medicine  as  a  germicide.  As  such  it  has 
the  distinct  advantage  that,  after  it  has  acted,  only  water  re- 
mains, which  is  harmless. 

The  usual  3  per  cent  solution  on  the  market  is  also  called 
dioxogen ;  it  frequently  is  diluted  further  as  required.  It  is 
kept  in  brown  bottles,  in  a  cool  place,  and  is  generally  very 
slightly  acidified,  which  greatly  reduces  the  rate  of  its  decom- 
position by  neutralizing  the  alkali  that  is  taken  up  from  the 
glass  of  the  bottle. 

Hydrogen  peroxide  is  manufactured  on  a  large  scale,  and 
most  of  it  is  employed  as  a  bleaching  agent.  Thus,  delicate 
silks,  ostrich  feathers,  ivory,  hair,  and  sponges  are  bleached 
with  hydrogen  peroxide.  It  is  used  to  change  dark-colored 
living  hair  to  lighter  color.  It  is  also  employed  similarly  in 
changing  the  color  of  furs.  In  these  bleaching  processes,  hy- 
drogen peroxide  is  used  because  it  is  a  mild  agent,  which  does 
not  injure  these  animal  tissues  as  much  as  other  bleaching 
agents  do.  Hydrogen  peroxide  is  also  used  in  photography 
to  remove  the  last  traces  of  sodium  thiosulphate  from  the  pho- 
tographic plates,  after  the  latter  have  been  "fixed."  In  ana- 
lytical chemistry  it  is  frequently  employed  as  an  oxidizing  agent. 

Ozonic  Acid.  —  Baeyer  and  Villiger  have  described  an  oxide 
of  hydrogen  which  contains  still  more  oxygen  than  hydrogen 
peroxide.  This  compound,  to  which  the  formula  HO2  or  H2O4 
has  been  assigned,  has  been  called  ozonic  acid,  because  it  may 
be  regarded  as  formed  by  the  addition  of  ozone  to  water  :  — 

O3  +  H20  =  H204. 

Ozonic  acid  has  not  yet  been  isolated,  but  Baeyer  and  Villiger 
regard  the  peroxide  of  potassium  K2O4,  for  instance,  as  a  salt 
of  ozonic  acid,  the  two  potassium  atoms  having  replaced  the 
hydrogen  atoms. 


CHAPTER  VIII 

THE   HALOGENS 

The  Halogen  Family.  —  The  elements  that  belong  to  this 
group  are  fluorine,  chlorine,  bromine,  and  iodine.  Of  these 
chlorine  is  the  most  common  and  the  most  abundant  in  nature. 
Its  properties  have  already  been  discussed.  Fluorine,  bromine, 
and  iodine  form  with  hydrogen  the  compounds  hydrogen 
fluoride  or  hydrofluoric  acid  HF,  hydrogen  bromide  or  hydro- 
bromic  acid  HBr,  and  hydrogen  iodide  or  hydriodic  acid  HI. 
These  compounds  are  analogous  to  hydrogen  chloride  or  hydro- 
chloric acid  HC1.  By  replacing  the  hydrogen  of  these  hydro- 
halogen  acids  by  means  of  sodium,  the  sodium  salts,  sodium 
fluoride  NaF,  sodium  chloride  NaCl,  sodium  bromide  NaBr,  and 
sodium  iodide  Nal  are  formed.  These  salts  are  quite  similar 
to  one  another ;  and  as  common  salt  is  a  member  of  the  group, 
the  elements  fluorine,  chlorine,  bromine,  and  iodine  have  been 
termed  the  halogens,  meaning  salt  formers.  This  must  not 
bp  taken  to  mean  that  all  salts  contain  one  of  these  four  ele- 
ments, for  such  is  not  at  all  the  case. 

With  the  exception  of  fluorine,  the  halogens  unite  with 
oxygen  and  hydrogen  to  form  certain  acids.  Chlorine  and 
iodine  also  unite  with  oxygen  to  form  oxides.  Furthermore, 
the  halogens  form  compounds  with  one  another,  with  the  metals, 
and  with  many  other  elements. 

We  shall  now  take  up  the  compounds  which  chlorine  forms 
with  oxygen  and  hydrogen,  after  which  the  remaining  halogens 
and  their  principal  compounds  will  be  considered. 

Compounds  of  Chlorine  with  Oxygen.  — There  are  three  of  these 
compounds,  namely,  chlorine  monoxide  C12O,  chlorine  dioxide 
C1O2,  and  chlorine  heptoxide  C12O7.  These  are  all  very  unstable 
substances,  decomposing  readily  into  chlorine  and  oxygen.  They 
are  not  formed  by  direct  interaction  of  chlorine  and  oxygen. 

Chlorine  monoxide  is  formed  when  chlorine  acts  on  cold  mer- 
curic oxide  :  — 

2  HgO  +  2  C12  =  HgO  •  HgCl2  +  C12O. 
96 


THE   HALOGENS  97 

It  is  a  brownish  yellow  gas,  which  may  be  condensed  to  a  liquid 
boiling  at  +5°.  The  substance,  especially  when  liquefied,  is 
highly  explosive.  It  detonates  when  heated  or  subjected  to 
concussions ;  but  in  the  sunlight  it  soon  decomposes  into 
chlorine  and  oxygen  without  explosion. 

Chlorine  dioxide  is  formed  when  potassium  chlorate  is  treated 
with  concentrated  sulphuric  acid.  The  reaction  may  be  re- 
garded as  taking  place  in  two  steps,  thus :  — 

(1)  KC1O3  -f  H2SO4  =  KHSO4  -f  HC1O3. 

chloric  acid 

(2)  3  HC103  =  HC104  +  H20  +  2  C1O2. 

perchloric 
acid 

Chlorine  dioxide  is  also  called  chlorine  peroxide.  It  is  a  yellow 
gas  which  may  be  condensed  to  a  liquid,  boiling  at  +9.9°. 
Solid  chlorine  dioxide  melts  at  —  79°.  The  substance  is  very 
explosive.  Its  odor  resembles  that  of  chlorine.  In  the  sunlight 
it  slowly  decomposes  into  the  elements.  It  is  a  powerful 
oxidizing  agent.  Sugar  mixed  with  potassium  chlorate  bursts 
into  flame  when  touched  with  a  drop  of  concentrated  sulphuric 
acid  ;  for  thus  chlorine  peroxide  is  liberated,  which  at  once 
attacks  the  sugar  violently.  Phosphorus  introduced  into 
chlorine  peroxide  gas  at  once  takes  fire.  When  the  gas  is 
touched  with  a  red-hot  iron,  it  explodes. 

Chlorine  heptoxide  is  formed  by  the  action  of  phosphorus 
pentoxide  on  perchloric  acid.  The  action  simply  consists  of 
the  elimination  of  a  molecule  of  water  from  two  molecules  of 
perchloric  acid :  — 

2  HC104  =  H2O  +  C1207. 

Chlorine  heptoxide  is  a  colorless  oil  which  boils  at  82°.  On 
percussion  it  explodes  with  violence,  also  when  brought  in  con- 
tact with  a  flame.  It  is  therefore  a  dangerous  substance  to 
handle,  and  great  care  must  be  exercised  in  distilling  it. 

Hypochlorous  Acid  and  Hypochlorites.  —  When  chlorine  mon- 
oxide acts  on  water  a  solution  of  hypochlorous  acid  is  formed:  — 

C12O  +  H2O  =  2  HOC1. 

Hypochlorous  acid  is  known  only  in  solution  and  in  form  of  its 

salts. 


98  OUTLINES   OF  CHEMISTRY 

When  caustic  potash  solution  is  treated  with  chlorine  at 
room  temperatures,  the  following  change  occurs :  — 

2  KOH  +  Cla  =  KOC1  +  KC1  +  H2O. 

potassium        potassium 
hypochlorite        chloride 

A  perfectly  analogous  change  occurs  when  chlorine  acts  on  cal- 
cium hydroxide,  slaked  lime  :  — 

2  Ca(OH)2  +  2  C12  =  Ca(OCl)2  +  CaCl2  +  2  H2O. 

calcium 
hypochlorite 

The  product  is  bleaching  powder  or  so-called  chloride  of  lime. 

It  consists  of  calcium  hypochlorite  Ca(OCl)2  and  calcium 
chloride  CaCl2.  The  formula  of  bleaching  powder  is,  however, 

best  expressed  thus :  Ca</  Q,  ,  for  the  substance  really  con- 
tains no  calcium  chloride,  since  it  lacks  the  hygroscopicity  of 
the  latter  salt.  Furthermore,  alcohol  will  not  extract  calcium 
chloride  from  bleaching  powder,  though  calcium  chloride  is 
soluble  in  alcohol. 

By  treating  calcium  hypochlorite  with  very  dilute,  cold  nitric 
acid  HNO3,  hypochlorous  acid  is  liberated,  thus:  — 

Ca(OCl)2  +  2  HN03  =  Ca(NO3)2  +  2  HOC1. 

Hypochlorous  acid  readily  loses  oxygen  and  passes  over  into 
hydrochloric  acid,  especially  in  the  sunlight :  — 

HOC1  =  HC1  +  O. 

The  nascent  oxygen  thus  liberated  readily  oxidizes  substances 
like  coloring  matters,  and  hence  hypochlorous  acid  is  an  oxidiz- 
ing and  bleaching  agent.  Hypochlorous  acid  has  twice  the 
bleaching  power  possessed  by  chlorine  water  containing  the 
same  amount  of  chlorine,  as  is  evident  from  the  following 
equations :  — 

2  HOC1  =  2  HC1  +  02. 
2C12  +  2H20  =  4HCH-02. 

Hypochlorous  acid  readily  decomposes  in  sunlight  into  oxy- 
gen and  hydrochloric  acid.  Concentrated  solutions  readily 
form  chloric  acid  and  hydrochloric  acid  :  — 

3  HC1O  =  2  HC1  +  HC1O3. 


THE   HALOGENS  99 

HypoMorites  are  also  oxidizing  agents.  They  part  with  their 
oxygen  and  pass  over  into  chlorides.  Thus  calcium  hypochlo- 
rite  slowly  forms  calcium  chloride  and  oxygen  :  — 

Ca(OCl)2  =  CaCl2  +  Oa. 

Upon  this  fact  depends  the  bleaching  action  of  bleaching  powder, 
and  also  its  disinfecting  action,  for  the  oxygen  liberated  destroys 
organic  matter.  Javelle  water  is  made  by  treating  bleaching 
powder  with  sodium  carbonate.  By  treating  bleaching  powder 
with  sulphuric  acid  all  the  chlorine  is  liberated,  thus  :  — 


The  chlorine  liberated  then  acts  upon  water,  forming  hydro- 
chloric acid  and  oxygen,  the  latter  destroying  the  coloring 
matter  to  be  bleached.  Hence,  in  using  bleaching  powder  in 
practice  it  is  generally  treated  with  an  acid. 

Chloric  Acid  and  Chlorates.  —  When  a  solution  of  potassium 
hypochlorite  is  heated,  the  following  change  occurs  :  — 

3  KC1O  =  2  KC1  +  KC1O3. 

potassium 
chlorate 

Potassium  chlorate  may  be  formed  directly  by  saturating  a  hot 
solution  of  caustic  potash  with  chlorine,  thus  :  — 

6  KOH  +  3  C12  =  5  KC1  +  KC1O3  +  3  H2O. 

As  potassium  chlorate  KC1O3  is  much  less  soluble  in  water 
than  potassium  chloride  KC1,  the  former  readily  crystallizes 
from  a  hot  saturated  solution  on  cooling. 

By  treating  potassium  chlorate  with  dilute  sulphuric  acid, 
chloric  acid,  whose  composition  is  represented  by  the  formula 
HC1O3,  is  liberated:  — 

2  KClOg  +  H2SO4  =  K2SO4  +  2  HC1O3. 

This  reaction  is  perfectly  analogous  to  that  of  making  hydro- 
chloric acid  :  — 

2  KC1  +  H2SO4  =  K2SO4  +  2  HC1. 

Chloric  acid  is  known  only  in  solution  and  in  form  of  its  salts. 
Its  anhydride  C12O5  is  not  known  at  all.  Chloric  acid  solutions, 
forming  thick,  colorless  sirups  of  specific  gravity  1.25,  have  been 
obtained.  They  contain  40  per  cent  of  the  free  acid  and  corre- 
spond approximately  to  the  formula,  HC1O3  4-  7  H2O. 


100  OUTLINES  OF   CHEMISTRY 

Attempts  to  concentrate  the  solution  further  always  result  in 
decomposition  of  the  chloric  acid  into  chlorine,  oxygen,  and 
perchloric  acid.  The  sirupy  solution  of  chloric  acid  oxidizes 
linen,  wood,  paper,  and  other  organic  material  very  rapidly, 
with  evolution  of  light  and  heat.  The  aqueous  solutions  of  the 
acid  are  much  more  stable  than  those  of  hypochlorous  acid ; 
still,  on  standing  perchloric  acid  is  formed  in  them,  especially 
in  the  sunlight.  The  salts  of  chloric  acid,  namely  the  chlorates, 
are  much  more  stable  than  the  hypo chlorite s. 

Perchloric  Acid  and  Perchlorates.  —  When  potassium  chlorate 
is  melted,  it  gives  off  oxygen  slowly  and  then  becomes  nearly 
solid,  forming  potassium  chloride  and  potassium  perchlorate  :  — 

4KC103  =  3KC104+KC1. 

The  potassium  perchlorate  KC1O4  is  much  less  soluble  in  water 
than  potassium  chloride,  hence  the  latter  -salt  may  readily  be 
separated  from  the  former.  Sodium  perchlorate  NaClO4  is 
found  in  small  amounts  in  Chili  saltpeter. 

On  heating  potassium  perchlorate,  it  gives  up  all  of  its 
oxygen,  passing  over  into  potassium  chloride.  Hence,  the 
formation  of  potassium  perchlorate  is  really  an  intermediate 
step  in  the  making  of  oxygen  by  heating  potassium  chlorate. 

By  treating  potassium  perchlorate  with  strong  sulphuric 
acid,  perchloric  acid  is  formed :  — 

KC104  +  H2S04  =  KHS04  +  HC1O4. 

Perchloric  acid  HC1O4  may  also  be  produced  by  heating  chloric 
acid,  or  by  exposing  the  latter  to  sunlight :  — 

3  HC1O3  =  C12  +  2  O2  +  H2O  +  HC1O4. 

Perchloric  acid,  prepared  by  carefully  distilling  a  mixture  of 
potassium  perchlorate  and  sulphuric  acid  in  a  partial  vacuum, 
is  a  colorless,  very  corrosive  liquid  which  fumes  strongly  in 
the  air.  It  has  a  specific  gravity  of  1.782  at  15.5°  and  a 
boiling  point  of  about  40°  at  60  mm.  pressure.  It  is  the  most 
stable  of  the  oxy-acids  of  chlorine  ;  still,  it  cannot  be  kept  long 
even  in  the  dark,  for  after  a  few  days  decomposition  with  violent 
explosion  occurs.  The  acid  is  a  dangerous  product.  In  con- 
tact with  the  skin  it  produces  wounds  that  are  painful  and 
very  slow  to  heal.  A  few  drops  put  on  paper,  wood,  etc. 
causes  these  substances  to  burst  into  flames,  while  a  drop  of 


THE   HALOGENS  101 

the  acid  on  charcoal  produces  a  violent  explosion.  These 
phenomena  occur  because  perchloric  acid  is  very  rich  in  oxy- 
gen, with  which  it  parts  readily,  thus  producing  violent  oxida- 
tion accompanied  with  sudden  liberation  of  much  heat. 

The  anhydride  of  perchloric  acid  is  chlorine  heptoxide  C12O7, 
which,  as  has  been  stated,  is  produced  by  abstracting  water 
from  perchloric  acid  by  treatment  with  phosphorus  pentoxide. 

Nomenclature  and  General  Relations. — The  following  table 
presents  the  formulae  and  names  of  the  compounds  of  chlorine 
with  oxygen  and  hydrogen  :  — 

HC1,  hydrochloric  acid.  KC1,  potassium  chloride. 

HC1O,  hypochlor0MS  acid.  KC1O,  potassium  hypochlon'fe. 

(HC1O2,  chlorous  acid).  KC1O2,  potassium  chlorite. 

HC1O3,  chloric  acid.  KC1O3,  potassium  chlorate. 

HC1O4,  perchloric  acid.  KC1O4,  potassium  perchlorate. 

Chlorous  acid  HC1O2  is  not  known  in  the  free  state ;  but  its 
salts,  like  potassium  chlorite  KC1O2,  are  known.  The  latter,  for 
instance,  is  formed  together  with  potassium  chlorate  when  chlo- 
rine dioxide  acts  on  caustic  potash  :  — 

2  KOH  +  2  C1O2  =  KClp2  4-  KC1O3  +  H2O. 

The  above  table  presents  an  interesting  series  of  compounds. 
Beginning  with  hydrochloric  acid  and  its  salt  potassium  chlo- 
ride, each  member  of  the  series  contains  one  atom  of  oxygen 
more  than  the  preceding.  Hydrochloric  acid  is  a  very  stable 
compound ;  but  hypochlorous  acid  is  very  unstable.  On  the 
other  hand,  chloric  acid  is  more  stable  than  hypochlorous  acid, 
and  perchloric  acid  is  the  most  stable  of  the  three  known  oxy- 
acids  of  chlorine.  The  salts  of  these  acids,  obtained  by  replac- 
ing the  hydrogen  of  the  acid  by  a  metal,  are  much  more  stable 
than  the  corresponding  acids.  Such  salts  form  articles  of  com- 
merce. Their  uses  will  be  considered  more  fully  later. 

The  names  given  the  oxy-acids  of  chlorine  and  their  corre- 
sponding salts  afford  an  excellent  illustration  of  the  system  of 
naming  a  series  of  acids  of  increasing  oxygen  content  and  the 
salts  which  they  form.  From  the  table  it  appears  that  HC1O3 
is  called  chloric  acid  and  its  salts  chlorates ;  the  acid  which  is 
richest  in  oxygen,  HC1O4,  is  called  perchloric  acid,  and  its  salts 
perchlorates ;  the  acid  containing  less  oxygen  than  chloric  acid, 


102  OUTLINES  OF   CHEMISTRY 

namely  HC1O2,  is  termed  chlorous  acid  and  its  salts  the  chlorto  ; 
whereas  the  acid  containing  still  less  oxygen,  HC1O,  is  called 
hypochlorous  acid  and  its  salts  hypochlorites.  Finally,  HC1, 
which  contains  no  oxygen  at  all,  is  termed  hydrochloric  acid, 
which  distinguishes  it  sufficiently  from  HC1O3,  chloric  acid. 
This  method  of  naming  acids  and  their  corresponding  salts  is 
generally  applied  in  chemistry  whenever  a  similar  series  of  com- 
pounds is  found.  The  ic  acid  forms  the  ate  salt ;  the  ous  acid 
forms  the  ite  salt ;  the  "hypo  .  .  ous  acid  forms  the  hypo  . .  ite  salt ; 
and  the  per  .  .  ic  acid  forms  the  per  . .  ate  salt.  Numerous  other 
illustrations  of  this  will  be  met  in  our  further  considerations. 

Occurrence,  Preparation,  and  Properties  of  Fluorine.  —  This 
element  is  widely  distributed  in  nature.  It  occurs  in  large 
quantities,  but  always  in  combination  with  other  elements.  It 
is  chiefly  found  combined  with  calcium  as  fluorspar,  calcium 
fluoride  CaF2,  which  crystallizes  in  octahedra  and  in  cubes  like 
common  salt.  In  Greenland,  fluorine  is  found  in  the  mineral 
cryolite,  which  is  a  fluoride  of  sodium  and  aluminum,  the  com- 
position of  which  is  expressed  by  the  formula  (NaF)3 .  A1F3. 
In  many  minerals  and  siliceous  rocks  fluorine  occurs  in  small 
quantities,  in  combination  with  calcium  and  other  metals. 
Fluorides  also  are  found  in  small  quantities  in  sea  water,  in 
many  mineral  waters,  in  the  ashes  of  plants,  and  in  the  teeth 
and  the  bones  of  animals.  Fluorspar  has  been  known  for  a  very 
long  time.  It  melts  at  red  heat,  and  has  been  used  as  a  flux 
in  metallurgical  processes  as  early  as  the  fifteenth  century.  It 
used  to  be  called  fluate  of  lime.  The  name  "  fluorine  "  comes 
from  the  use  of  fluorspar  as  a  flux. 

Fluorine  was  not  isolated  till  1886,  when  Henri  Moissan  pre- 
pared it  by  passing  the  electric  current  through  dry,  liquid 
hydrofluoric  acid  HF,  in  which  potassium  hydrogen  fluoride 
KHF2  had  been  dissolved,  in  order  to  have  the  liquid  conduct 
electricity.  The  solution  was  placed  in  a  tube  made  of  platinum 
(Fig.  29),  the  stoppers  being  made  of  fluorspar.  The  electrodes 
were  made  of  an  alloy  of  platinum  and  iridium.  The  apparatus 
was  kept  at  —  23°  C.,  and  the  fluorine  was  collected  in  a  platinum 
tube,  the  ends  of  which  were  closed  with  transparent  plates  of 
fluorspar.  The  difficulty  in  isolating  fluorine  lies  in  the  fact 
that  the  element  combines  so  readily  with  other  elements. 
Moissan  found  later  that  perfectly  pure  fluorine  attacks  glass 


THE  HALOGENS 


103 


but  very  slowly  indeed,  so  that  the  gas  may  be  collected  in  glass 
vessels.  It  has  also  been  demonstrated  that  a  copper  vessel 
may  be  used  instead  of 
one  of  platinum  in  pre- 
paring fluorine. 

Fluorine  is  a  gas  of 
a  light,  greenish  yellow 
color  and  a  strong  pun- 
gent odor.  It  may  be 
condensed  to  a  liquid 
which  boils  at  — 187°. 
By  chilling  the  liquid 
with  liquid  hydrogen, 
it  freezes,  the  white 
crystals  formed  melting 
at  -  223°.  Fluorine 
gas  is  19  times  as  heavy 
as  hydrogen.  Its  molec-  FlG  2g. 

ular  weight  is  therefore 

38 ;  and  since  its  atomic  weight  is  19.0,  the  formula  of  fluorine 
is  F2.  The  atomic  weight  of  fluorine  has  been  determined 
from  the  analysis  of  calcium  fluoride. 

Fluorine  is  the  most  active   of  all  the  elements.     It  acts  on 
water,  yielding  ozone  and  hydrofluoric  acid  :  — 

3  H20  +  3  F2  =  6  HF  +  O3. 

It  unites  with  hydrogen  with  great  violence  in  the  dark  at 
ordinary  temperatures,  and  even  at  —  253°  solid  fluorine  still 
acts  with  explosive  violence  on  liquid  hydrogen,  according  to 
Dewar  and  Moissan.  Most  of  the  non-metallic  elements  unite 
directly  with  fluorine  at  ordinary  temperatures  with  evolution 
of  heat  and  light.  Iron,  lead,  barium,  strontium,  calcium,  so- 
dium, and  potassium  are  acted  upon  by  fluorine  at  ordinary 
temperatures;  magnesium,  aluminum,  manganese,  nickel,  and 
silver  burn  in  fluorine  when  slightly  heated.  At  ordinary 
temperatures  gold  and  platinum  are  not  attacked,  but  between 
300°  and  400°  they  are  converted  into  fluorides.  Copper  is 
acted  upon  at  ordinary  temperatures,  a  coating  of  cuprous 
fluoride  being  formed  at  once  on  the  metal,  which  is  thus  pro- 
tected from  further  action.  Oxygen,  chlorine,  nitrogen,  and 


104  OUTLINES   OF   CHEMISTRY 

argon  do  not  unite  with  fluorine.  Organic  substances  generally 
burn  in  fluorine  gas.  Hydrochloric  acid  gas  is  decomposed  by 
fluorine  with  explosive  violence  :  — 

2  HC1  +  F2  =  2  HF  +  C12. 

Dry  glass  is  but  very  slowly  attacked  by  fluorine,  but  in  pres- 
ence of  hydrofluoric  acid  or  water,  even  in  traces,  glass  is 
rapidly  destroyed. 

Hydrofluoric  Acid.  When  calcium  fluoride  is  treated  with 
sulphuric  acid,  hydrofluoric  acid  is  formed  :  — 

CaF2  +  H2SO4  =  CaSO4  +  2  HF. 

The  experiment  is  carried  on  in  a  platinum  or  lead  dish,  for 
hydrofluoric  acid  acts  upon  glass  or  porcelain.  The  process  of 
making  the  acid  is  perfectly  analogous  to  that  of  preparing 
hydrochloric  acid.  To  obtain  hydrofluoric  acid  which  is  anhy- 
drous, i.e.  free  from  water,  potassium  hydrogen  fluoride  KHFa 
is  heated  to  redness  in  a  platinum  retort :  — 

KHF2=KF  +  HF. 

Hydrofluoric  acid  is  a  liquid  whose  boiling  point  is  + 19.4°. 
Solid  hydrofluoric  acid  melts  at  —  92.3°.  When  perfectly  dry, 
the  liquid  does  not  act  upon  glass.  In  presence  of  moisture, 
however,  glass  is  rapidly  attacked,  fluorides  and  water  being 
formed.  Glass  consists  essentially  of  the  silicates  of  sodium 
and  calcium,  Na2SiO3  and  CaSiO3.  When  hydrofluoric  acid 
acts  upon  these,  the  following  changes  occur :  — 

CaSiO3  +  6  HF  =  CaF2  +  SiF4  +  3  H2O. 
Na2SiO3  +  6  HF  =  2  NaF  +  SiF4  +  3  H2O. 

The  compound  silicon  tetrafluoride  SiF4  is  a  gas,  and  so  es- 
capes. Calcium  fluoride  is  soluble  in  acids,  so  that  when  glass 
is  attacked  by  hydrofluoric  acid,  it  is  dissolved.  Use  is  made 
of  this  fact  in  the  chemical  analysis  of  glass  and  other  silicates, 
also  in  etching  glass.  In  the  latter  process  the  glass  is  first 
coated  with  paraffin ;  the  design  is  traced  in  the  paraffin  coat- 
ing so  as  to  expose  the  portions  of  the  glass  to  be  etched,  and 
the  whole  is  then  treated  either  with  the  fumes  of  hydrofluoric 
acid  or  with  an  aqueous  solution  of  the  latter.  When  the  par- 
affin is  finally  removed,  the  design  is  found  etched  into  the  sur- 


THE   HALOGENS  105 

face  of  the  glass.     This  process  is  used  in  marking  graduations 
on  glass  utensils,  thermometers,  etc. 

Because  hydrofluoric  acid  attacks  glass,  it  is  kept  in  rubber  or 
wax  bottles.  It  is  very  soluble  in  water,  and  fumes  in  contact 
with  moist  air.  The  concentrated  solution  boils  at  120°,  and 
contains  about  36  to  38  per  cent  of  the  anhydrous  acid. 

Hydrofluoric  acid  is  a  dangerous  substance,  for  it  is  very 
poisonous.  When  inhaled  it  produces  death.  In  contact  with 
the  skin  it  produces  swellings,  pains,  and  wounds  that  are  very 
slow  to  heal. 

At  100°  hydrogen  fluoride  is  about  ten  times  as  heavy  as 
hydrogen  ;  this  leads  to  the  molecular  formula  HF.  But  at  25° 
the  vapor  of  hydrofluoric  acid  is  nearly  20  times  as  heavy  as 
hydrogen,  which  leads  to  the  formula  H2F2.  The  acid  is  prone 
to  form  acid  salts  like  KHF2  and  NaHF2.  The  other  hydro- 
halogens  do  not  form  analogous  compounds. 

Occurrence,  Preparation,  and  Properties  of  Bromine.  —  Like 
chlorine,  bromine  does  not  occur  in  nature  except  in  combina- 
tion with  other  elements.  Bromine  is  generally  found  in 
nature  with  chlorine  in  salt  deposits.  And  just  as  chlorine 
occurs  mainly  in  form  of  sodium  chloride,  so  bromine  occurs 
chiefly  as  sodium  bromide.  Bromine  is  widely  distributed  in 
nature,  but  it  is  not  found  anywhere  in  very  large  quantities. 
In  sea  water,  sodium  bromide  and  magnesium  bromide  are 
found.  Together  these  constitute  from  0.3  to  1.3  per  cent  of 
the  residue  obtained  by  evaporating  the  water.  In  the  Stass- 
furt  salt  beds,  bromine  occurs  as  magnesium  bromide.  The 
salt  wells  of  West  Virginia,  Ohio,  and  Michigan  furnish  most 
of  the  bromine  used  in  the  United  States.  Here  the  element 
occurs  as  sodium  bromide  together  with  common  salt.  On 
evaporating  the  brine  the  sodium  chloride  is  first  deposited,  it 
being  less  soluble  than  sodium  bromide.  From  the  mother  liquor 
sodium  bromide,  together  with  some  common  salt,  is  obtained 
by  further  evaporation.  In  the  year  1910  the  United  States 
produced  245,437  pounds  of  bromine  valued  at  $41,684. 

Bromine  was  discovered  in  1826  by  Balard,  who  prepared  it 
from  the  residue  obtained  by  evaporating  sea  water. 

The  method  of  preparing  bromine  is  the  same  as  that  of  pre- 
paring chlorine,  thus :  — 
2  NaBr  +  3  HaSO4  +  MnOa  =  2  NaHSO4+MnSO4  +  2 


106  OUTLINES   OF   CHEMISTRY 

Chlorine  will  readily  replace  bromine,  and  so  this  method  may 
be  employed  in  making  bromine  :  — 

2  NaBr  +  Cla  =  2  NaCl  +  Br2. 
MgBr2+Cl2=MgCl2  +  Br2. 

This  method  is  used  in  manufacturing  bromine  in  Michigan 
and  at  Stassfurt. 

Bromine  is  the  only  non-metallic  element  which  is  a  liquid 
at  ordinary  temperatures.  It  is  dark  reddish  brown  in  color, 
boils  at  59°,  and  has  a  specific  gravity  of  3.188  at  0°.  At  -  7.5° 
it  freezes  to  a  dark  brown  solid,  and  at  —98°  it  crystallizes 
from  carbon  disulphide  in  carmine  red  needles.  At  room  tem- 
peratures bromine  vaporizes  readily.  It  irritates  the  eyes  and 
the  mucous  membranes  of  the  mouth  and  throat  and  has  an 
extremely  disagreeable  odor  ;  whence  its  name  bromine,  mean- 
ing a  stench.  In  contact  with  the  skin  it  produces  wounds 
that  are  painful  and  difficult  to  heal. 

Bromine  dissolves  in  water.  The  solution  has  the  color  of 
bromine  and  is  known  as  bromine  water.  At  room  temperature 
20°  the  saturated  solution  contains  about  3  per  cent  bromine. 
On  cooling  the  solution  to  about  0°  a  hydrate  of  the  composi- 
tion Br2  +  10  H2O  separates  out.  This,  however,  readily 
decomposes  when  warmed  to  room  temperature. 

In  its  chemical  behavior,  bromine  closely  resembles  chlorine. 
With  metals  and  a  large  number  of  other  elements  it  unites 
directly,  forming  bromides.  Thus  arsenic  and  antimony  will 
burn  in  bromine,  which  also  reacts  vigorously  with  phosphorus 
and  sulphur.  On  the  other  hand,  it  does  not  unite  with  carbon 
or  oxygen  directly.  It  acts  violently  on  potassium  ;  but  dry 
sodium  may  even  be  heated  with  bromine  up  to  200°  before 
appreciable  action  begins.  Bromine  turns  starch  paste  yellow. 
It  bleaches  like  chlorine,  only  much  more  slowly.  The  bleach- 
ing action  depends  upon  the  fact  that  bromine,  like  chlorine,  acts 
upon  water,  liberating  oxygen,  which  attacks  organic  coloring 
matters,  thus :  — 

H20  +  Br2  =  2  HBr  +  O. 

The  atomic  weight  of  bromine  is  79.92,  and  as  its  vapor  is 
about  79.5  times  heavier  than  hydrogen,  its  molecular  weight 
is  159.84,  and  its  molecular  formula  is  Bra. 


THE  HALOGENS  107 

Hydrobromic  Acid.  —  Hydrogen  bromide,  or  hydrobromic  acid 
HBr,  may  be  formed  by  direct  union  of  hydrogen  with  bromine, 
which  occurs  when  hydrogen  charged  with  bromine  vapor  is 
ignited.  By  treating  sodium  bromide  with  sulphuric  acid 
hydrogen  bromide  is  formed,  just  as  hydrogen  chloride  forms 
when  common  salt  is  similarly  treated,  thus :  — 

NaBr  +  H2SO4  =  NaHSO4  +  HBr. 

However,  in  this  case  a  portion  of  the  hydrobromic  acid  liber- 
ated at  once  reacts  with  some  of  the  sulphuric  acid,  forming 
bromine,  water,  and  sulphur  dioxide.  That  is  to  say,  some  of 
the  hydrobromic  acid  reduces  sulphuric  acid  :  — 

H2S04  +  2  HBr  =  SO2  +  2  H2O  +  Br2. 

Thus,  pure  hydrobromic  acid  cannot  be  obtained  by  treating 
sodium  bromide  with  sulphuric  acid,  for  the  product  contains 
free  bromine  and  also  sulphur  dioxide. 

Pure  hydrogen  bromide  is  formed  when  phosphorus  tribro- 
mide  PBr3  or  phosphorus  pentabromide  PBr5  is  acted  upon 
by  water,  thus  :  — 

PBr3  +  3  H20  =  H3P03  +  3  HBr. 

phosphorous 
acid 

PBr6  +  4  H20  =  H3PO4  +  5  HBr. 

phosphoric 
acid 

Phosphorous  and  phosphoric  acids  are  not  volatile,  but  hydro- 
bromic acid  is,  and  so  the  latter  can  readily  be  separated  from 
the  former. 

The  apparatus  used  for  making  hydrobromic  acid  is  shown 
in  Fig.  30.  The  flask  F  contains  red  phosphorus  covered 
with  a  little  water.  Bromine  is  gradually  added  by  opening 
the  cock  0.  Phosphorus  bromide  is  formed,  which  is  at  once 
decomposed  by  the'  water,  yielding  hydrobromic  acid.  The 
latter  generally  contains  some  bromine  vapor,  which  is  removed 
by  allowing  the  gas  to  pass  over  pumice  covered  with  moist  red 
phosphorus  in  the  U-tube  A. 

Hydrobromic  acid  is  a  colorless  gas  of  strong  pungent  odor. 
It  may  be  condensed  to  a  liquid  which  boils  at  —  64.9°  under 
738.2  mm.  pressure.  It  forms  colorless  crystals  which  melt  at 
—  88°.  It  furnes  strongly  in  the  air,  and  is  very  soluble  in 


108 


OUTLINES   OF   CHEMISTRY 


water,  one  volume  of  the  latter  absorbing  about  600  volumes  of 
hydrobromic  acid  gas  at  10°. 

Hydrobromic  acid  is  a  strong  acid  which  readily  attacks 
many  metals,  forming  bromides  of  the  metals  and  liberating 
hydrogen,  thus :  — 

Mg  +  2  HBr  =  MgBr2  +  H2. 
Zn  +  2HBr=ZnBr2  +  H2. 

In  general,  the  chemical  behavior  is  like  that  of  hydrochloric 
acid.  Like  the  chlorides  of  the  metals,  the  bromides  are  gen- 
erally soluble  in  water ;  and  just  as  the  chloride  of  silver  AgCl 


FIG.  30. 


is  insoluble,  so  the  bromide  of  silver  AgBr  is  also  insoluble. 
Further,  the  bromide  of  lead  PbBr2  and  mercurous  bromide 
HgBr  are  difficultly  soluble  like  the  corresponding  chlorides. 

When  hydrobromic  acid  is  treated  with  chlorine,  hydrochloric 
acid  and  bromine  are  produced  :  — 

2HBr  +  Cl2=2HCl  +  Br2. 

On  boiling,  a  strong  aqueous  solution  of  hydrobromic  acid 
becomes  weaker,  and  a  weak  solution  becomes  stronger,  till 
finally  a  solution  containing  from  47.4  to  47.8  per  cent  of 
hydrobromic  acid  is  formed.  This  solution  then  distills  over 
unchanged  in  concentration  at  752  to  762  mm.  pressure.  How- 
ever, by  distilling  it  at  other  pressures  its  strength  is  changed. 


THE  HALOGENS  109 

Thus  at  16  mm.  pressure  the  acid  that  distills  over  contains 
51.6  per  cent  HBr. 

Hydrobromic  acid  is  40.45  times  heavier  than  hydrogen. 
Its  molecular  weight  is  therefore  80.9.  By  weight  it  contains 
1.008  grams  of  hydrogen  to  every  79.92  grams  of  bromine. 
From  these  data,  its  formula  is  HBr.  When  dry  hydrobromic 
acid  gas  is  treated  with  metallic  sodium  in  an  apparatus  like 
that  used  in  investigating  the  composition  of  HC1  (Fig.  23), 
it  is  found  that  the  hydrogen  liberated  occupies  one  half  of  the 
volume  of  the  hydrobromic  acid  taken. 

At  very  high  temperatures  hydrobromic  acid  gas  decomposes 
into  hydrogen  and  bromine.  This  reaction,  which  is  a  rever- 
sible one,  like  the  decomposition  of  hydrochloric  acid  gas  and 
of  water  in  the  gaseous  state  at  very  high  temperatures,  is 
another  typical  case  of  dissociation,  and  may  be  represented 
thus  :  — 


The  term  dissociation  is  only  applied  to  reversible  reactions  in 
which  a  compound  is  decomposed  into  products  which  may 
again  unite  to  form  the  original  compound  as  the  pressure, 
temperature,  or  amount  of  material  contained  in  unit  of  volume 
is  varied.  We  shall  have  occasion  to  refer  to  other  instances  of 
dissociation. 

Oxy-acids  of  Bromine.  —  There  are  no  oxides  of  bromine 
known,  and  but  two  oxy-acids  have  thus  far  been  prepared. 
They  are  hypobromous  acid  HBrO  and  bromic  acid  HBrO3, 
the  latter  being  known  in  aqueous  solution  only. 

Hypobromous  acid  and  its  salts,  the  hypobromites,  are  prepared 
in  a  manner  analogous  to  the  preparation  of  hypochlorous  acid 
and  hypochlorites.  Thus,  by  action  of  bromine  water  on  mer- 
curic oxide,  hypobromous  acid  HBrO  results,  just  as  hypochlo- 
rous acid  is  formed'when  chlorine  water  acts  on  mercuric  oxide. 
The  changes  are  expressed  as  follows  :  — 

2  Bra  +  H2O  +  HgO  =  HgBr2  +  2  HBrO. 
2  Cl?  +  H20  +  HgO  =  HgCl2  +  2  HC1O. 

While  hypochlorous  acid  is  known  only  in  aqueous  solutions, 
hypobromous  acid  may  be  isolated  by  distillation  in  a  partial 
vacuum  at  40°.  The  solution  of  the  acid  in  water  is  straw-yel- 
low in  color.  The  acid  readily  decomposes  into  hydrobromic 


110  OUTLINES  OF   CHEMISTRY 

acid  and  oxygen,  and  is,  therefore,  like  hypochlorous  acid,  a 
strong  oxidizing  and  bleaching  agent. 

When  bromine  acts  on  a  cold  solution  of  caustic  alkali,  Tiypo- 
bromites  are  formed.  The  process  is  analogous  to  the  formation 
of  hypochlorites,  thus :  — 

2  KOH  +  Br2  =  KBr  +  KBrO  +  H2O. 

Hypobromites,  like  hypochlorites,  are  unstable  compound's,  readily 
giving  up  oxygen.  The  hypobromite  solutions  yield  bromates 
readily,  especially  at  higher  temperatures  :  — 

3KBrO  =  2KBr  +  KBr03; 

or  in  warm  caustic  potash  solution,  bromine  at  once  forms 
potassium  bromate,  thus :  — 

3  Br2  +  6 KOH  =  5  KBr  +  3  H2O  +  KBrO3. 

Bromic  Acid  and  Bromates. — When  silver  bromate  AgBrO3 
is  treated  with  bromine  and  water,  bromic  acid  is  formed :  — 

5  AgBrO3  +  3  Br2  +  3  H2O  =  5  AgBr  +  6  HBrO3. 

It  is  also  formed  when  dilute  sulphuric  acid  acts  on  barium 
bromate :  — 

Ba(BrO3)2  +  H2SO4  =  BaSO4  +  2  HBrO8 ; 
or  when  chlorine  is  passed  into  bromine  water,  thus :  — 
Br2  +  6  H2O  +  5  C12  =  10  HC1  +  2  HBrO3. 

Bromic  acid  is  very  similar  to  chloric  acid  in  its  behavior.  At 
100°  the  aqueous  solution  decomposes,  yielding  oxygen  and 
bromine.  The  pure  anhydrous  acid  has  not  been  prepared. 

On  heating  potassium  bromate,  it  yields  ^oxygen  and  potas- 
sium bromide,  without,  however,  first  forming  a  potassium  per- 
bromate.  In  this  respect  the  behavior  of  potassium  bromate 
KBrO3  differs  from  that  of  potassium  chlorate. 

By  melting  potassium  bromide  with  potassium  chlorate,  po- 
tassium bromate  results  :  — 

KC1O3  +  KBr  =  KBrO3  +  KC1. 

We  thus  see  that  under  these  conditions  the  bromate  is  more 
stable  than  the  chlorate,  and  the  chloride  more  stable  than  the 
bromide. 


THE   HALOGENS  111 

Uses  of  Bromine  and  its  Compounds.  —  Bromine  is  used  in  the 
manufacture  of  ctyestuffs  from  coal-tar  products.  In  medicine 
potassium  bromide  is  used  as  a  sedative.  In  photography  silver 
bromide  is  used  in  the  sensitized  plates. 

History  and  Occurrence  of  Iodine.  —  This  element  is  a  solid  at 
room  temperatures.  It  was  discovered  in  1812  by  Courtois, 
who  evolved  it  from  the  ashes  of  seaweeds.  It  forms  beautiful 
violet  vapors,  whence  its  name  iodine,  meaning  violet-colored. 
Indeed  it  was  the  color  of  the  vapor  that  led  to  the  discovery 
of  iodine.  The  substance  was  then  examined  by  Sir  Humphry 
Davy  and  by  Gay-Lussac  ;  the  latter  in  1815  established  its 
elementary  character. 

Like  bromine,  iodine  always  occurs  in  nature  associated  with 
chlorine.  It  has  been  reported  by  Wanklyn  that  the  water 
from  the  spring,  Woodhall  Spa,  near  Lincoln,  Nebraska,  con- 
tains iodine  in  minute  quantities ;  but  with  this  singular  excep- 
tion, iodine  has  always  been  found  in  combination  with  other 
elements,  chiefly  with  sodium,  potassium,  magnesium,  and  cal- 
cium in  form  of  iodides  and  iodates.  In  sea  water  it  occurs  in 
extremely  minute  quantity.  Seaweeds,  particularly  those  grow- 
ing in  deeper  waters,  like  the  genera  Fucus  and  Laminaria, 
assimilate  iodine  and  store  it  up  in  their  bodies.  The  ashes  of 
such  seaweeds  are  termed  kelp  in  Scotland  and  varech  in  Nor- 
mandy, and  from  these  iodine  is  prepared.  However,  the 
chief  source  of  iodine  at  present  is  the  crude  Chili  saltpeter, 
or  caliche  NaNO3,  in  which  iodine  occurs  mainly  as  sodium 
iodate  NaIO3.  The  amount  of  iodine  in  caliche  is,  however, 
only  about  0.2  per  cent.  Besides  occurring  in  seaweeds, 
iodine  is  found  in  many  sponges,  oysters,  and  other  sea  animals, 
in  cod-liver  oil,  in  some  fresh-water  plants,  in  coal,  and  in 
the  thyroid  glands  of  animals.  In  combination  with  silver, 
copper,  and  lead  it  occurs  as  iodides,  though  these  minerals 
are  rare.  Many  mineral  springs  contain  minute  amounts  of 
iodine,  and  in  deposits  of  common  salt  the  element  generally 
occurs  in  small  quantity.  Thus  it  is  evident  that  iodine  is 
quite  widely  distributed  in  nature,  though  it  is  nowhere  present 
in  large  amounts. 

Preparation  of  Iodine.  —  From  the  ashes  of  seaweeds,  iodine 
is  liberated  by  treatment  with  sulphuric  acid  and  manganese 
dioxide,  or  by  passing  chlorine  through  the  solution,  which  con- 


112 


OUTLINES   OF   CHEMISTRY 


tains  the  iodine  mainly  in  the  form  of  sodium  iodide.      The 
equations  expressing  the  changes  that  occur  are  as  follows  :  — 


(2) 


2  Nal  +  C12  =  2  NaCl  +  T2. 


It  will  be  observed  that  equation  (1)  is  perfectly  analogous  to  the 
process  of  making  chlorine  or  bromine  from  chlorides  or  bro- 
mides by  treatment  with  sulphuric  acid  and  manganese  dioxide. 
Further,  equation  (2)  is  analogous  to  the  process  of  making 
bromine  from  a  bromide  by  treatment  with  chlorine.  On  the 
coasts  of  France  and  Scotland  the  seaweeds  are  gathered,  dried, 
and  burned,  the  latter  process  being  carried  on  in  closed  retorts 

so  that  no  iodine  is  lost  by 
volatilization.  The  charcoal 
remaining  after  the  ash  has 
been  leached  out  of  it  is  sim- 
ilar to  animal  charcoal.  It 
readily  absorbs  odors  and  is 
used  as  a  deodorant.  The 
iodine  is  then  liberated  from 
the  solution  by  means  of  one 
of  the  processes  just  men- 
tioned. It  is  purified  by 
volatilizing  it  and  condens- 
ing the  vapor,  which  forms 
crystals.  The  process  of 
FlG  31  vaporizing  a  solid  without 

melting    it    and    condensing 

the  vapor  to  the  solid  state  is  called  sublimation.  To  get  pure 
iodine  the  latter  is  mixed  with  potassium  iodide,  and  the  mix- 
ture is  heated  so  as  to  volatilize  the  iodine,  which  is  condensed 
on  cool  surfaces  in  form  of  crystals.  In  this  way  bromine  and 
chlorine  remain  behind,  in  combination  with  potassium.  Figure 
31  shows  a  simple  laboratory  apparatus  for  subliming  iodine, 
and  Fig.  32  represents  an  arrangement  for  resubliming  raw 
iodine  on  a  commercial  scale. 

From  Chili  saltpeter,  in  which  iodine  occurs  as  sodium  iodate 
together  with  smaller  amounts  of  sodium  iodide  and  magnesium 
iodide,  iodine  is  prepared  by  treatment  with  sodium  bisulphite. 
The  reaction  which  takes  place  is  as  follows  :  — 


THE   HALOGENS 


113 


FIG.  32. 


2  NaIO3  +  5  NaHSO3  =  2  Na2SO4  +  3  NaHSO4  +  H2O  +  I2. 
The  iodine  is  thus  obtained  in  precipitated  form  from  the  aque- 
ous solution.  It  is  allowed  to  settle  and  is  then  collected  and 
purified  by  sublimation.  The  quantity  of  iodine  produced 
annually  from  Chili 
saltpeter  is  about 
800  tons,  which  is 
somewhat  more  than 
half  of  the  total  pro- 
duction. Of  recent 
years,  the  process  of 
obtaining  iodine 
from  seaweeds  has 
been  improved,  the 
best  method  consist- 
ing of  lixiviating 
the  seaweeds  without 
previous  charring. 
In  this  way  much 
less  iodine  is  lost, 
and  the  remains  of 

the  seaweeds  are  worked  up  into  cHgin,  which  is  like  gelatine. 
Thus  the  method  of  preparing  iodine  from  seaweeds  has  again 
become  profitable. 

Properties  of  Iodine.  —  Iodine  is  a  grayish  black,  lustrous 
solid  which  crystallizes  in  plates  that  belong  to  the  rhombic 
system.  From  solutions  in  alcohol  or  hydriodic  acid  beautiful 
crystals  may  be  obtained.  Iodine  really  has  a  metallic  luster. 
It  is  brittle  and  may  be  pulverized  readily.  Its  specific  gravity 
is  4.95  at  17°.  It  melts  at  116.1°,  forming  a  reddish  brown 
liquid  which  boils  at  184°.  Iodine  volatilizes  perceptibly, 
though  slowly,  at  room  temperatures.  Its  vapors  are  violet- 
colored,  but  when  dense  they  appear  very  dark  and  opaque. 
Its  odor  reminds  one  of  that  of  chlorine  and  bromine,  but  it  is 
much  less  intense.  It  colors  the  skin  brown  and  exerts  an 
irritating  and  corrosive  action  upon  it. 

In  water  it  is  but  very  slightly  soluble,  about  1  part  in  5000. 

However,  water  containing  potassium  iodide  or  hydriodic  acid 

readily  dissolves  iodine.     These  solutions  are  brown,  as  is  also 

the   solution   of  iodine  in  alcohol,  which  is  called  tincture  of 

i 


114  OUTLINES   OF  CHEMISTRY 

iodine.  Iodine  furthermore  dissolves  in  hydrocarbon  oils,  chlo- 
roform, and  carbon  disulphide.  With  the  latter  it  forms  beau- 
tiful violet  solutions,  which  fact  is  frequently  used  in  detecting 
iodine  in  chemical  analysis.  We  have  already  learned  that 
iodine  turns  starch  paste  blue.  This  is  used  as  a  test  for  iodine 
and  also  for  starch  in  analytical  chemistry. 

Uses  of  Iodine.  —  The  solution  of  iodine  in  alcohol,  tincture 
of  iodine,  is  used  as  a  counter  irritant  in  medicine.  Iodine  is 
also  administered  internally  in  form  of  potassium  iodide  as  a 
specific  in  certain  diseases,  particularly  those  which,  like  goiter, 
are  caused  by  disturbances  in  the  thyroid  gland.  The  latter 
normally  contains  iodine  in  form  of  an  organic  compound 
known  as  thyroiodine.  This  also  occurs  in  the  thyroid  glands 
of  animals,  particularly  in  sheep,  from  which  source  it  is 
mainly  obtained.  It  is  administered  as  a  specific  for  goiter 
and  myxoedema. 

Iodine  is  also  used  in  the  manufacture  of  iodoform,  iodocrol, 
and  other  iodine  preparations.  These  are  used  principally  as 
antiseptics  in  healing  wounds.  In  synthetic  chemistry,  hydriodic 
acid  and  compounds  of  iodine  with  carbon  and  hydrogen  are 
frequently  employed. 

Hydriodic  Acid.  —  There  is  but  one  compound  of  hydrogen 
and  iodine  known  ;  namely,  hydriodic  acid  or  hydrogen  iodide. 
Its  composition  and  vapor  density  are  represented  by  the  for- 
mula HI.  It  may  be  prepared  by  passing  a  mixture  of 
hydrogen  and  iodine  vapor  through  a  red-hot  tube  containing 
platinum  in  a  finely  divided  state,  thus:  — 


The  reaction  is,  however,  incomplete  since  it  is  a  reversible  one, 
hydriodic  acid  decomposing  readily  into  iodine  and  hydrogen. 

By  treating  potassium  iodide  with  sulphuric  acid,  we  cannot 
obtain  hydriodic  acid  ;  for  the  latter  reduces  sulphuric  to  sul- 
phurous acid  far  more  readily  than  does  hydrobromic  acid:  — 

2  KI  +  3  H2S04  =  2  KHS04  +  SO2  +  2  H2O  +  2  I. 
However,  by  treating  potassium  iodide  with  hot,  concentrated 
phosphoric  acid,  hydriodic  acid  may  be  obtained,  thus  :  — 

KI  +  H3P04  =  KHaP04  +  HI. 

The  acid  is  best  prepared  by  decomposition  of  phosphorus 
iodide  by  water:  — 


THE    HALOGENS  115 

PT3  +  3  H20  =  HgPOg  +  3  HI ; 

or  by  simply  having  red  phosphorus  and  iodine  act  on  each 
other  in  presence  of  water.  In  this  way  phosphorus  iodide 
is  formed  and  then  decomposed  into  phosphorous  acid  and 
hydrogen  iodide:  — 

P  +  3  I  +  3  H20  =  H3P03  +  3  HI. 

Hydrogen  iodide  may  also  be  obtained  by  the  action  of  iodine 
on  hydrogen  sulphide  H2S  (which  see),  or  by  the  action  of 
hydrogen  sulphide  upon  cuprous  iodide  suspended  in  water, 

thus :  — 

H2S  +  I2  =  2HI  +  S. 
H2S  +  2CuI=2HI  +  Cu2S. 

These  methods  are  quite  similar  to  those  by  means  of  which 
pure  hydrobromic  acid  can  be  obtained.  Hydrochloric  acid 
may  also  be  prepared  by  similar  methods ;  but  it  is  not  at  all 
necessary  to  resort  to  these  in  this  case,  since  this  acid  is  much 
more  stable  than  hydrobromic  or  hydriodic  acid.  It  does  not 
reduce  sulphuric  acid,  and  can  therefore  readily  be  prepared 
by  the  action  of  the  latter  on  common  salt. 

At  room  temperatures  hydrogen  iodide  is  a  colorless  gas, 
which,  like  hydrochloric  and  hydrobromic  acids,  fumes  strongly 
in  the  air  and  is  very  soluble  in  water.  At  60°,  485  volumes 
of  hydriodic  acid  gas  are  absorbed  by  1  volume  of  water.  At 
0°  a  solution  of  specific  gravity  2.0  may  be  obtained  which 
contains  about  90  per  cent  hydrogen  iodide.  On  distillation, 
a  solution  of  hydrogen  iodide  behaves  like  the  corresponding 
solution  of  hydrogen  bromide  and  hydrogen  chloride.  On 
boiling,  a  concentrated  solution  becomes  weaker,  and  a  weak 
solution  becomes  more  concentrated,  till  finally  a  liquid  is 
obtained  which  contains  57  per  cent  HI.  This  boils  at  127° 
at  774  mm.  and  distills  over  without  change  of  composition. 
On  changing  the  pressure,  however,  the  composition  of  the  dis- 
tillate is  changed.  The  distillation  of  the  hydriodic  acid  must 
be  conducted  in  a  current  of  hydrogen  to  prevent  decomposi- 
tion of  the  acid. 

Pure  hydriodic  acid  may  be  condensed  to  a  colorless  liquid 
which  boils  at  —34.1°.  Solid  hydriodic  acid  forms  colorless 
crystals  which  melt  at  —50.8°.  The  vapor  of  hydriodic  acid 


116  OUTLINES  OF   CHEMISTRY 

is  62.92  times  heavier  than  hydrogen,  whence  its  molecular 
weight  is  126.84,  which  corresponds  fairly  well  to  the  formula 
HI;  for  the  atomic  weight  of  iodine  is  126.92,  and  the  calcu- 
lated molecular  weight  for  the  formula  HI  is  127.92. 

Hydriodic  acid  forms  iodides  and  hydrogen  when  treated 
with  many  metals.  These  salts  are  as  a  rule  soluble  in  water, 
the  exceptions  being  the  iodides  of  silver,  mercury,  and  lead. 
Hydriodic  acid  is  a  powerful  reducing  agent,  which  comes 
from  the  fact  that  it  readily  gives  up  its  hydrogen  to  oxidizing 
agents.  The  decomposition  of  hydriodic  acid  proceeds  more 
rapidly  in  the  light  and  at  higher  temperatures.  Its  reducing 
power  is  frequently  used  in  chemistry,  especially  in  the  investi- 
gation of  the  compounds  of  carbon. 

Oxide  of  Iodine.  —  But  one  oxide  of  iodine  is  known.  Its  com- 
position is  expressed  by  the  formula  I2O5.  It  is  the  anhydride 
of  iodic  acid,  and  is  prepared  by  heating  the  latter  to  170°, 
thus  :  — 

2HI03=H20  +  I206- 

When  the  oxide  is  dissolved  in  water,  the  acid  is  regenerated. 
Iodine  pentoxide  is  a  white  crystalline  solid,  which  decomposes 
into  its  elements  at  300°  ;  thus  it  is  much  more  stable  than  the 
oxides  of  chlorine. 

Oxy-acids  of  Iodine.  —  There  are  three  oxy-acids  of  iodine 
known,  namely  hypoiodous  acid  HIO,  iodic  acid  HIO3,  and 
periodic  acid  HIO4. 

A  dilute  solution  of  hypoiodous  acid  may  be  prepared  by 
shaking  together  mercuric  oxide,  iodine,  and  water:  — 

HgO  +  2  12  +  H20  =  2  HIO  +  HgI2. 

The  method  is  thus  similar  to  the  preparation  of  HC1O  and 
HBrO.  When  iodine  is  introduced  into  cold  caustic  alkali 
solutions,  a  colorless  liquid  harving  bleaching  power  results, 
due  to  the  formation  of  hypoiodites,  thus  :  — 

2  NaOH  +  12  =  NalO  +  Nal  +  H2O. 

Hypoiodites  are,  however,  extremely  unstable,  readily  passing  over 
into  iodates,  especially  on  warming,  thus  :  — 


Iodic  acid  HIO3  is  perfectly  analogous  to  chloric  and  bromic 


THE   HALOGENS  117 

acids.  It  is,  however,  much  more  stable  than  the  latter.  It 
may  be  formed  by  treating  barium  iodate  with  sulphuric  acid  :  — 

Ba(IO3)2  +  H2SO4  =  BaSO4  +  2  HIO3 ; 

or  by  oxidation  of  iodine  either  by  means  of  chlorine  or  nitric 
acid,  thus :  — 

1  +  3  H20  +  5  Cl  =  5  HC1  +  HIO8 

31  +  5  HN03  =  5  NO  +  H2O  +  3  HIO3. 

nitric  acid        nitric  oxide 

lodic  acid  readily  gives  up  oxygen  and  is  consequently  a  good 
oxidizing  agent.  Thus  in  contact  with  hydriodic  acid,  both 
acids  are  decomposed,  yielding  water  and  iodine  :  — 

HIO3  +  5  HI  =  3  H2O  +  3 12. 

The  salts  of  iodic  acid  are  called  the  iodates.  The  potassium 
and  sodium  salts  readily  dissolve  in  water ;  but,  in  general,  the 
salts  of  other  metals  are  sparingly  soluble.  On  being  heated, 
the  iodates  behave  like  the  bromates.  The  iodates  of  sodium 
and  potassium  yield  iodides  and  oxygen,  whereas  other  iodates 
decompose  into  oxides  of  the  metal,  iodine,  and  oxygen.  The 
iodates  of  potassium  and  sodium  readily  unite  with  one  or  two 
additional  molecules  of  iodic  acid,  forming  acid  salts.  Thus, 
KIO3  -  HIO8  and  KIO3  •  2  HIO3  are  known.  The  chlorates  and 

OO  DO 

bromates  do  not  thus  add  on  chloric  and  bromic  acid. 

Periodic  acid  is  formed  by  the  action  of  iodine  upon  an  aque- 
ous solution  of  perchloric  acid  :  — 

2  HC104  +  4  H20  +  I2  =  C12  +  2  (HIO4  -  2  H2O) . 

Periodic  acid  has  the  composition  corresponding  to  the  formula 
HIO4  •  2  H2O,  or  as  it  is  often  written,  H5IO6.  The  acid  of 
the  formula  HIO4  has  never  been  obtained.  Periodic  acid  forms 
colorless,  transparent,  deliquescent,  prismatic  crystals  that  melt 
at  133°,  while  at  140°  they  are  entirely  decomposed,  forming 
water,  oxygen,  and  iodine  pentoxide,  thus  :  — 

2H6T06  =  5H20  +  02  +  I206- 

Periodic  acid  is  a  strong  oxidizing  agent.  Periodates  are  gener- 
ally difficultly  soluble  in  water.  Sodium  periodate  may  readily 
be  prepared  by  the  interaction  of  chlorine,  sodium  hydroxide, 
and  sodium  iodate,  thus  :  — 

C12  +  3  NaOH  4-  NaIO3  =  2  NaCl  +  Na2H3IO6. 


118  OUTLINES   OF   CHEMISTRY 

By  heating  barium  iodate,  the  periodate  of  barium  may  be  ob- 
tained, thus  :  — 

5  Ba(I03)2  =  Ba5(I06)2  +  4  12  +  9  O2. 

It  will  thus  be  seen  that  while  periodic  acid  and  the  periodates 
are  analogous  to  perchloric  acid  and  the  perchlorates,  the  fact 
that  periodic  acid  has  the  composition  HIO4  •  2  H2O  or  H5IO6 
leads  to  a  more  complicated  series  of  salts  than  we  have  in  case 
of  the  perchlorates. 

Compounds  of  the  Halogens  with  Each  Other.  —  By  passing 
chlorine  over  iodine,  a  dark  reddish  brown  liquid  not  unlike 
bromine  in  appearance  is  formed.  It  is  very  volatile  and  has 
an  exceedingly  pungent  odor.  It  is  about  3  times  as  heavy  as 
water  and  boils  at  about  101°,  during  which  process  it  suffers 
partial  decomposition.  Its  composition  corresponds  to  the 
formula  IC1  ;  it  is  iodine  monochloride.  Two  modifications  of 
this  compound  have  been  described,  the  one  melting  at  24.7°, 
and  the  other  at  13.9°.  Iodine  monochloride  does  not  turn 
starch  paste  blue.  By  contact  with  water  it  is  decomposed  :  — 

3  H2O  +  5  IC1  =  HIO3+  5  HC1  +  2  I2. 

When  iodine  is  treated  with  chlorine  in  excess,  or  when 
iodine  monochloride  is  further  treated  with  chlorine,  yellow, 
needle-like  crystals  are  formed  having  the  composition  IC18. 
They  are  iodine  trichloride.  They  may  be  purified  by  sublima- 
tion at  ordinary  temperatures.  On  heating,  they  decompose 
into  chlorine  and  iodine  monochloride,  but  on  cooling  the  tri- 
chloride forms  again,  thus  :  — 


Water  dissolves  iodine  trichloride.  The  solution  has  great 
germicidal  power  and  is  consequently  used  as  an  antiseptic. 

With  bromine,  iodine  forms  a  crystalline  compound,  iodine 
monobromide,  of  the  composition  IBr.  It  has  chemical  proper- 
ties similar  to  those  of  iodine  monochloride.  It  melts  at  36°. 

A  compound  of  iodine  with  fluorine,  iodine  pentafluoride  IF5, 
is  also  known.  It  is  formed  by  direct  union  of  the  elements. 
It  is  a  colorless  liquid  which  boils  at  97°  and  solidifies  at  8°. 
On  heating  it  to  400°,  it  suffers  decomposition.  Water  decom- 
poses it  into  hydrofluoric  and  iodic  acids. 


THE   HALOGENS  119 

General  Relations  of  the  Halogens  to  One  Another.  —  Fluorine, 
chlorine,  bromine,  and  iodine  increase  in  atomic  weight  in  the 
order  named.  With  increasing  atomic  weight  their  melting 
points  and  boiling  points  rise,  their  specific  gravities  increase, 
and  their  color  becomes  more  intense.  Thus,  with  increasing 
atomic  weight,  we  have  here  an  increase  in  the  degree  of  con- 
densation of  matter,  as  it  were. 

While  the  physical  properties  thus  show  a  regular  change 
with  increasing  atomic  weight,  the  chemical  properties  also 
exhibit  regularity  of  change.  So  the  affinity  for  hydrogen  is 
greatest  in  the  case  of  fluorine,  and  least  in  the  case  of  iodine. 
The  general  chemical  activity  of  the  halogens  diminishes  as  the 
atomic  weight  increases.  So  fluorine  is  by  far  the  most  active 
element  of  the  group,  and  iodine  the  least  active.  However, 
for  oxygen  iodine  has  a  much  greater  affinity  than  fluorine, 
which  unites  with  oxygen  neither  directly  nor  indirectly.  In- 
deed, in  case  of  the  oxygen  compounds  of  the  halogens  the  sta- 
bility increases  with  the  atomic  weight  of  the  halogen,  being 
greatest  in  the  iodine  compounds. 

It  is  interesting  to  note  that  the  atomic  weight  of  bromine, 
79.92,  is  approximately  equal  to  one  half  the  sum  of  the  atomic 
weights  of  chlorine,  35.46,  and  iodine,  126.92.  We  shall  meet 
more  such  groups  of  three  elements  in  which  a  similar  relation 
holds.  The  consideration  of  the  relations  between  the  atomic 
weights  of  the  elements  and  their  physical  and  chemical  proper- 
ties has  led  to  a  classification  of  the  elements  known  as  the 
periodic  system,  which  will  be  considered  when  more  of  the 
elements  have  been  studied. 


CHAPTER   IX 

ACIDS,   BASES,    SALTS, 
HYDROLYSIS,  MASS  ACTION,  AND  CHEMICAL  EQUILIBRIUM 

Acids.  — In  connection  with  the  study  of  oxygen  it  was  found 
that  this  element  readily  unites  with  non-metals  like  phos- 
phorus, sulphur,  and  carbon,  forming  oxides  which,  when 
dissolved  in  water,  have  a  sour  taste  and  redden  blue  litmus. 
These  oxides  are^^guently  called  acidic  oxides  or  acid-form- 
ing oxides,  for  wiSRater  they  form  acids.  Thus  when  sulphur 
burns  in  oxygen  su||iur  dioxide  results:  — 

S  +  O2=SO2. 

When  conducted  into  water,  sulphur  dioxide  unites  with  the 
water,  forming  sulphurous  acid  H2SO3,  thus:  — 

S02  +  H20  =  H2S03. 

It  is  possible  to  form  a  similar  acid  of  higher  oxygen  content. 
By  passing  sulphur  dioxide  mixed  with  oxygen  over  red-hot, 
finely  divided  platinum,  a  higher  oxide  of  sulphur,  namely 
sulphur  trioxide  SO3,  is  formed.  This  is  a  white  crystalline 
solid  which  greedily  unites  with  water,  forming  sulphuric  acid 
H2SO4.  The  changes  may  be  expressed  as  follows  :  — 

2S02  +  02  =2  S03. 
SO3  +  H2O=H2SO4. 

Similarly,  when  phosphorus  is  burned  in  oxygen  phosphorus 
pentoxide  P2O5  is  formed,  which  readily  unites  with  water, 
forming  phosphoric  acid  H3PO4  :  — 

P205  +  3H20  =  2H3P04. 

Again,  when  carbon  is  burned  in  oxygen,  carbon  dioxide  CO2 
is  produced,  which,  when  dissolved  in  water,  forms  carbonic  acid 
H2CO3,  thus :  — 

C02  +  H20  =  H2C03. 

Carbonic  acid  H2CO3  has  not  been  isolated  ;  it  exists  merely  in 
solution.  The  combination  of  water  and  carbon  dioxide  is  very 

120 


ACIDS,    BASES,   SALTS,   CHEMICAL  EQUILIBRIUM  121 

weak,  and  carbonic  acid  is  but  slightly  sour  to  the  taste,  red- 
dens blue  litmus  slowly,  and  acts  in  all  respects  much  more 
feebly  than  the  other  acids  just  mentioned.  It  is  a  good  exam- 
ple of  a  weak  acid. 

It  will  be  recalled  that  the  fact  that  oxides  like  the  above 
yield  sour  or  acidic  substances  originally  led  to  the  idea  that  it 
is  the  oxygen  that  imparts  these  acidic  characteristics  to  the  com- 
pounds. Indeed,  oxygen  received  its  name  in  accordance  with 
this  notion.  However,  we  have  seen  that  the.  halogens  form  a 
series,  of  compounds  with  hydrogen,  namely,  HF,  HC1,  HBr,  and 
HI,  which  are  all  pronounced  acids,  in  that  they  are  sour  to  the 
taste,  redden  litmus,  and  attack  metals  like  magnesium,  zinc, 
and  iron,  evolving  hydrogen  and  yielding  compounds  consisting 
of  the  halogen  and  the  metal  employed.  These  latter  products 
are  salts  of  the  metal. 

When  chlorine  was  discovered,  it  was  looked  upon  as  an 
oxide  (as  oxidized  hydrochloric  acid) ;  for  it  was  well  known 
that  chlorine  is  an  acid-forming  substance,  and  consequently  it 
was  thought  that  it  must  contain  oxygen,  which  was  regarded  as 
the  essential  element  in  every  acid.  In  fact,  it  was  not  till  iodine 
was  discovered  that  the  correct  view  of  chlorine  as  an  element 
was  really  established.  For  this  reason  the  discovery  of  iodine 
and  the  proof  that  it  is  elementary  in  character  was  of  great 
importance  in  the  development  of  the  idea  of  the  real  nature  of 
an  acid.  Thus,  from  the  study  of  the  hydrohalogens,  which 
are  all  pronounced  acids,  came  the  true  notion  that  hydrogen, 
and  not  oxygen,  is  the  essential  constituent  of  every  acid.  An 
acid  is  a  compound  containing  hydrogen  which  may  be  replaced 
by  a  metal,  the  product  formed  being  a  salt.  Acids  commonly 
have  sour  taste  and  redden  blue  litmus ;  but  we  shall  learn  of 
acids  that  are  so  weak  that  they  do  neither  of  these  things ;  yet 
they  are  nevertheless  acids,  because  they  contain  hydrogen 
which  may  be  replaced  by  a  metal,  the  product  formed  being  a 
salt.  While  it  is  thus  true  that  there  are  very  pronounced 
acids  that  contain  no  oxygen,  still  it  must  be  stated  that  after 
all  by  far  the  great  majority  of  acids  do  contain  oxygen,  and 
often  it  is  by  the  addition  of  the  latter  element  that  the  acidic 
character  is  produced. 

Since  phosphoric  acid  H3PO4,  sulphuric  acid  H2SO4,  sulphur- 
ous acid  H2SO3,  and  carbonic  acid  H2CO3,  may  be  formed  by 


122  OUTLINES  OF  CHEMISTRY 

the  addition  of  water  to  the  oxides  P2O5,  SO3,  SO2,  and  C02, 
respectively,  in  accordance  with  the  equations  already  given 
above,  these  oxides  are  called  the  anhydrides  of  the  respective 
acids. 

When  an  acid  acts  on  a  metal  like  zinc  or  magnesium,  hydro- 
gen is  evolved  and  a  salt  is  formed,  thus  :  — 

Zn  -+  H2SO4  =  ZnSO4  +  H2. 
Mg  +  2  HC1  =  MgCl2  -f  H2. 

A  salt  is  thus  one  of  the  products  of  the  interaction  of  an  acid 
and  a  metal.  It  is  possible  to  form  salts  by  other  means,  how- 
ever, as  will  be  shown  below. 

Bases.  —  We  have  seen  in  connection  with  the  study  of 
hydrogen  that  when  this  element  is  liberated  by  the  action  of 
sodium  or  potassium  on  water,  caustic  soda  or  caustic  potash 
results,  thus  :  — 

H2O  +  Na  =  NaOH  +  H. 

H2O  +  K  =  KOH  +  H. 

The  solutions  of  sodium  hydroxide  and  potassium  hydroxide 
are  alkaline  in  character.  They  turn  red  litmus  blue,  and  when 
treated  with  an  acid  they  become  neutral ;  that  is,  they  do  not 
affect  either  red  or  blue  litmus.  When  these  hydroxides  are 
treated  with  an  acid,  they  are  said  to  be  neutralized.  In  this 
process  the  acid  is,  of  course,  also  neutralized.  The  interaction 
of  an  alkaline  hydroxide  with  an  acid  is  a  mutual  act,  resulting  in 
the  neutralization  of  both  compounds.  On  evaporating  the  neutral 
solution  it  is  found  that  a  salt  has  been  formed.  The  reaction 
or  neutralizing  sodium  hydroxide  with  hydrochloric  acid  may 
be  expressed  thus  :  — 

NaOH  +  HC1  =  NaCl  +  H2O. 

When,  for  example,  potassium  hydroxide  is  neutralized  with 
sulphuric  acid,  the  following  change  takes  place  :  — 

2  KOH  +  H2SO4  =  K2SO4  +  2  H2O 

Hydroxides  of  metals  which  are  thus  capable  of  reacting  with 
acids,  forming  salts  and  water,  are  called  basic  hydroxides  or 
bases.  Hence  a  base  is  an  hydroxide  or  oxide  of  a  metal  which 
will  react  with  an  acid,  forming  (1)  a  neutral  substance  called  a 
salt,  and  (2)  water. 


ACIDS,   BASES,    SALTS,    CHEMICAL   EQUILIBRIUM  123 

Elements  which  are  thus  capable  of  uniting  with  the  hydroxyl 
radical  OH  to  form  bases  are  called  base-forming  elements,  while 
those  that  form  acids  by  union  with  hydrogen  are  called  acid- 
forming  elements.  Hydroxides  of  some  of  the  elements,  how- 
ever, are  capable  of  acting  as  bases  toward  more  acidic  hydrox- 
ides, and  as  acids  toward  hydroxides  that  are  more  basic  than 
themselves.  This  will  be  more  evident  as  we  proceed  in  our 
considerations. 

Salts.  —  From  what  has  been  stated  in  connection  with  the 
consideration  of  acids  and  bases,  the  nature  of  salts  is  already 
sufficiently  characterized.  Thus,  a  salt  is  a  neutral  compound 
resulting  as  a  product  of  the  interaction  of  an  acid  with  a  base  ; 
or  a  salt  is  a  neutral  compound  which  is  formed  when  the  hydrogen 
of  an  acid  is  replaced  by  a  metal.  So  it  appears  that  a  salt  may 
be  formed  in  the  following  ways  :  — 

(1)  By  the  neutralization  of  an  acid  with  a  base,  as,  for  exam- 
ple, sodium  sulphate  Na2SO4  is  formed  when  sodium  hydroxide 
and  sulphuric  acid  act  on  each  other  :  — 

2  NaOH  +  H2SO4  =  Na2SO4  +  2  H2O. 

(2)  By  the  action  of  a  metal  on  an  acid,  thus  :  — 

2  Na  4-  H2SO4  =  Na2SO4  +  H2. 

.(3)  It  is  possible,  however,  to  form  a  salt  by  the  direct  action 
of  two  oxides,  a  base-forming  oxide  or  basic  oxide,  and  an  acid- 
forming  oxide  or  acidic  oxide,  on  each  other,  thus :  — 

Na20  +  S03  =  Na2SO4. 

(4)  It  is  also  possible  to  form  a  salt  by  direct  union  of  a  base- 
forming  element  with  an  acid-forming  element,  thus  :  — 

Na  +  Cl  =  NaCl. 

In  all  cases,  however,  one  may  think  of  the  salt  as  derived 
from  some  acid  whose  hydrogen  has  been  replaced  by  the  metal. 
So,  though  in  (3)  the  sulphate  of  sodium  was  made  by  the 
union  of  sodium  oxide  and  sulphur  trioxide,  that  is  sulphuric 
anhydride,  one  may  think  of  the  product  Na2SO4  as  derived 
from  H2SO4  in  which  the  hydrogen  is  replaced  by  sodium. 
Likewise,  all  sulphates  may  be  regarded  as  similarly  derived  from 
sulphuric  acid.  The  latter  may  in  turn  be  looked  upon  as 
hydrogen  sulphate,  i.e.  a  salt  in  which  hydrogen  is  the  basic 


124  OUTLINES   OF   CHEMISTRY 

element;  indeed,  any  add  may  be  regarded  as  a  salt  of  hydrogen. 
Thus,  hydrochloric  acid  HC1  is  the  chloride  of  hydrogen. 
Sodium  chloride  NaCl,  whether  made  by  the  direct  union  of 
the  elements  as  in  (4)  above,  or  by  the  action  of  sodium 
hydroxide  upon  hydrogen  chloride,  may  yet  be  regarded  as 
derived  from  HC1  in  which  the  hydrogen  is  replaced  by  sodium. 

Similarly,  all  fluorides  may  be  regarded  as  derived  from  HF, 
the  hydrogen  of  which  has  been  replaced  by  the  metal.  All  iodides 
may  similarly  be  considered  as  derived  from  HI,  all  chlorates  from 
HC1O3,  etc. 

Older  View  of  the  Process  of  Salt  Formation.  —  Salts  were 
formerly  considered  as  the  result  of  the  union  of  a  basic  oxide 
with  an  acidic  oxide.  Thus,  sodium  sulphate  was  regarded  as 
sodium  oxide  Na2O  plus  sulphur  trioxide  SO3,  and  the  formula 
of  the  salt  was  written  Na2O  •  SO3.  Similarly,  calcium  carbon- 
ate was  regarded  as  made  up  of  calcium  oxide  CaO,  plus  carbon 
dioxide  CO2,  and  the  formula  of  calcium  carbonate  was  conse- 
quently CaO  •  CO2.  Again,  ferrous  sulphate  was  considered  as 
ferrous  oxide  FeO  plus  sulphur  trioxide,  x  thus  :  FeO  •  SO3. 
This  was  the  dualistic  way  of  writing  which  was  in  vogue  dur- 
ing the  former  half  of  the  last  century,  and  it  is  not  to  be  denied 
that  it  had  many  advantages.  So  these  formula  indicated  at 
once  that  the  salts  can  be  formed  by  direct  union  of  the  acidic 
and  basic  oxides  ;  and  since  it  is  true  that  many  of  these  salts 
when  strongly  heated  decompose  into  the  basic  and  acidic 
oxides,  the  formulse  also  in  a  simple  way  represented  this  fact. 
For  instance,  on  heating,  calcium  carbonate  yields  calcium  oxide 
and  carbon  dioxide  :  — 

CaCO3=CaO  +  CO2; 
and  ferrous  sulphate  decomposes  thus  :  —  • 


According  to  the  older  view  the  process  of  solution  of  a 
metal  like  zinc  or  iron  in  sulphuric  acid  consisted  of  two  steps. 
First,  when  the  zinc  was  introduced  into  dilute  sulphuric  acid, 
the  metal  was  oxidized  to  zinc  oxide  ZnO,  hydrogen  being 
simultaneously  liberated  from  the  water;  and  second,  zinc 
oxide  would  then  combine  with  the  sulphur  trioxide,  forming 
zinc  sulphate  ZnO  •  SO3.  Sulphuric  acid  was  regarded  as  SO3 
dissolved  in  water.  This  dualistic  way  of  writing  the  formulse 


ACIDS,   BASES,    SALTS,    CHEMICAL  EQUILIBRIUM  125 

of  salts  was  strongly  defended  by  Berzelius,  and  it  was  only 
through  the  development  of  the  study  of  the  compounds  of 
carbon  and  of  electrochemistry  that  the  present  method  of  ex- 
pressing the  formulae  was  finally  adopted.  In  the  study  of 
some  of  the  complicated  silicates,  however,  the  old  way  of 
writing  is  still  frequently  employed  with  distinct  advantages, 
as  will  be  seen  when  the  compounds  of  silicon  are  discussed. 
The  dualistic  formulae  of  Berzelius  were,  moreover,  also  based 
upon  electrochemical  ideas. 

Acid-  and  Base-forming  Elements.  —  In  general,  the  acid- 
forming  elements  are  the  non-metals,  and  the  base-forming  ele- 
ments are  the  metals.  Oxygen,  sulphur,  nitrogen,  phosphorus, 
carbon,  the  halogens,  etc.,  are  acid-forming  elements ;  and 
potassium,  sodium,  magnesium,  zinc,  lead,  copper,  etc.,  are  base- 
forming  elements.  However,  as  stated  above,  an  acidic  element 
may  act  as  a  basic  element  toward  a  still  more  acidic  element ; 
and  a  basic  element  may  act  as  an  acidic  element  toward  a  still 
more  basic  element.  Of  this  we  have  already  had  illustrations. 
Thus  while  zinc  acts  as  a  base  in  zinc  sulphate  ZnSO4,  in 
which  compound  sulphur  and  oxygen  form  the  acid  radical 
SO4,  in  potassium  zincate  K2ZnO2,  formed  thus, — 

Zn(OH)2  +  2  KOH  =  K2ZnO2  +  2  H2O, 

zinc  is  a  part  of  the  acid  radical  ZnO2.  So  toward  the  acid 
group  SO4  zinc  acts  as  a  base,  while  toward  the  strongly  basic 
potassium  the  zinc  forms  a  part  of  the  acidic  group  ZnO2.  Again, 
in  sodium  iodide  Nal,  sodium  is  the  basic  and  iodine  the  acidic 
element ;  whereas  in  iodine  chloride  IC1,  the  iodine  plays  the 
role  of  base  toward  the  more  acidic  chlorine.  Further,  in  phos- 
phorus trichloride  PC13,  phosphorus  is  the  basic  and  chlorine  the 
acidic  element,  whereas  in  phosphates,  like  sodium  metaphos- 
phate  NaPO3,  phosphorus  plays  the  role  of  an  acidic  element. 
Additional  examples  will  readily  occur  to  the  reader,  and  many 
others  will  be  met  in  our  further  study. 

The  distinction  between  acid-  and  base-forming  elements  is 
thus  not  a  sharp  one  ;  nevertheless,  from  what  has  been  stated, 
the  difference  between  acids  and  bases  can  generally  be  made 
without  difficulty. 

Other  Views  of  Solutions  of  Acids,  Bases,  and  Salts.  —  In  recent 
years  the  attempt  has  been  made  to  define  acids,  bases,  and  salts 


126  OUTLINES   OF   CHEMISTRY 

on  the  basis  of  the  behavior  of  their  solutions  toward  the  elec- 
tric current.  In  this  attempt  the  boiling  and  freezing  points 
of  dilute  solutions  have  also  been  a  prime  consideration.  This 
study  has  further  led  to  another  way  of  regarding  the  act  of 
neutralization  and  the  resulting  salt  solutions.  A  considera- 
tion of  these  views  of  acids,  bases,  and  salts  in  solution  will  be 
taken  up  later  in  connection  with  the  subjects  of  solutions  and 
electrolysis. 

Basicity  of  Acids;  Acid  Salts.  —  An  acid  which  contains  one 
replaceable  hydrogen  atom  in  its  molecule  is  called  a  monobasic 
acid;  one  that  contains  two  replaceable  hydrogen  atoms  is  called 
a  dibasic  acid  ;  etc.  There  are  also  tribasic,  tetrabasic,  and  penta- 
basic  acids.  For  example,  HC1,  HBr,  HI,  HC1O,  HBrO,  HIO, 
HC1O3,  HBrO3,  HIO3,  HC1O4,  HNO3  (nitric  acid),  are  all 
monobasic  acids.  With  a  univalent  basic  element  like  potas- 
sium or  sodium,  for  instance,  each  of  these  monobasic  acids 
forms  but  one  salt,  thus  :  — 

KC1,  Nal,  KBrO,  KC1O4,  NaNO3. 

Sulphuric  acid  H2SO4  and  carbonic  acid  H2CO3  are  dibasic 
acids,  since  they  contain  two  atoms  of  replaceable  hydrogen 
per  molecule.  In  the  neutralization  of  a  dibasic  acid  the  opera- 
tion may  take  place  in  two  steps,  thus  :  — 


(1)  H2SO4 

(2)  NaHS04  +  NaOH  =  Na2SO4    +  H2O. 

(1)  H2C03     +  KOH   =  KHC03  +  H2O. 

(2)  KHCO3  +KOH  =K2CO3     +  H2O. 

The  salt  NaHSO4  still  contains  replaceable  hydrogen,  i.e.  acidic 
hydrogen.  It  is  consequently  an  acid  salt,  as  contrasted  with 
Na2SO4,  in  which  all  the  hydrogen  has  been  replaced.  The 
latter  salt  is  a  neutral  or  normal  salt.  Acid  sodium  sulphate 
NaHSO4  is  also  called  sodium  bisulphate,  for  it  contains  twice 
as  much  acid  radical  per  same  amount  of  sodium  as  does  the 
normal  salt  Na2SO4,  which  is  also  called  bisodium  sulphate,  as 
well  as  simply  sodium  sulphate.  Similarly,  KHCO3  is  potas- 
sium bicarbonate,  and  K2CO3  is  bipotassium  carbonate,  or  simply 
potassium  carbonate. 

Thus  an  acid  salt  is  one  which  still  contains  hydrogen  that  is 
replaceable  by  a  metal. 


ACIDS,   BASES,   SALTS,   CHEMICAL  EQUILIBRIUM          127 

The  ability  of  an  acid  to  form  acid  salts  shows  that  the  acid 
is  not  monobasic.  Thus,  for  instance,  the  fact  that  hydro- 
fluoric acid  forms  acid  salts  like  KHF2  and  NaHF2  would  argue 
in  favor  of  the  view  that  the  acid  is  dibasic  in  character  and 
has  the  formula  H2F2,  rather  than  the  simple  formula  HF. 

In  phosphoric  acid  H3PO4  and  periodic  acid  H5TO6  we  have 
an  example  of  a  tribasic  and  a  pentabasic  acid,  respectively. 
The  hydrogen  atoms  of  these  acids  can  be  replaced  step  by 
step,  thus  forming  a  series  of  salts  which  grow  less  and  less 
acid  in  character.  For  instance,  in  case  of  phosphoric  acid,  we 
may  form  the  three  salts  KH2PO4,  K2HPO4,  and  K3PO4,  which 
are  called  primary,  secondary,  and  tertiary  potassium  phosphate, 
respectively.  We  may  also  call  these  salts  monopotassium  phos- 
phate, dipotassium  phosphate,  and  tripotassium  phosphate.  The 
commonest  salt  of  the  three  is  dipotassium  phosphate,  and  so  it 
is  generally  referred  to  simply  as  potassium  phosphate.  In  the 
case  of  the  pentabasic  periodic  acid,  we  have  a  still  greater 
range  of  possibility  of  formation  of  acid  salts ;  indeed,  the 
greater  the  number  of  replaceable  hydrogen  atoms  the  molecule  of 
an  acid  contains,  i.e.  the  greater  its  so-called  basicity,  the  larger 
is  the  number  of  acid  salts  that  it  is  able  to  form. 

Acidity  of  Bases.  — A  base  like  KOH  or  NaOH,  which  is  an 
hydroxide  of  a  univalent  metal,  is  called  a  monoacid  base,  for  it 
is  capable  of  reacting  with  an  acid,  and  thus  forming  one  mole- 
cule of  water  and  replacing  one  atom  of  acidic  hydrogen.  A 
base  like  Ca(OH)2  is  called  a  diacid  base,  for  it  is  capable  of 
reacting  with  an  acid,  forming  two  molecules  of  water  and  a  salt 
in  which  the  bivalent  metal  replaces  two  atoms  of  acidic  hydro- 
gen. Similarly,  we  may  have  triacid,  tetraacid,  and  pentaacid 
bases  like  antimonous  hydroxide  Sb(OH)3,  stannic  hydroxide 
Sn(OH)4,  and  antimonic  hydroxide  Sb(OH)5.  The  so-called 
acidity  of  a  base  consequently  is  determined  by  the  valence  of  the 
metal  of  the  base  ;  or  what  really  comes  to  the  same  thing,  by  the 
number  of  basic  hydroxyl  groups  the  molecule  of  the  base  contains. 

Basic  Salts.  —  Just  as  in  the  case  of  acids  that  contain  more 
than  one  hydrogen  atom  to  the  molecule  it  is  possible  to  secure 
acid  salts  by  replacing  these  hydrogen  atoms  step  by  step,  so 
in  the  case  of  bases  that  contain  more  than  one  basic  OH  group 
it  is  possible  to  neutralize  those  groups  step  by  step  by  means 
of  an  acid,  thus  forming  a  series  of  basic  salts.  For  example, 


128  OUTLINES   OF   CHEMISTRY 

(1)  Sb(OH)3  +  HC1  =  Sb(OH)2Cl  +  H20. 

(2)  Sb(OH)2Cl  +  HCl  =  Sb(OH)Cl2  +  H2O. 

(3)  Sb(OH)Cl2  +  HCl  =  SbCl3  +  H20. 

Thus  Sb(OH)2Cl  and  Sb(OH)Cl2  are  basic  salts,  for  they  still 
contain  an  excess  of  base  that  is  not  yet  neutralized.  The 
compound  Sb(OH)2Cl  readily  splits  off  water,  thus :  — 

Sb(OH)2Cl  =  SbOCl  +  H20. 

The  salt  SbOCl  is  called  antimony  oxy chloride;  it  is  also 
clearly  a  basic  salt,  for  it  is  still  capable  of  further  neutraliza- 
tion with  an  acid.  So  with  hydrochloric  acid  it  undergoes 
the  following  change:  — 

SbOCl  +  2  HC1  =  SbCl3  +  H2O. 

A.  basic  salt  is  one  that  contains  an  excess  of  base,  which  may  still 
be  neutralized  with  an  acid.  Basic  salts  are  frequently  met 
with,  and  they  often  split  off  water,  as  was  exemplified  above. 

Normal  Salts.  —  Normal  salts  are  those  that  contain  neither 
hydrogen  that  is  replaceable  by  a  metal  nor  an  excess  of  base  that 
may  still  be  neutralized  by  an  acid.  Thus  NaCl,  K2SO4,  CaCO3, 
are  normal  salts  of  hydrochloric,  sulphuric,  and  carbonic  acids, 
respectively.  From  what  has  been  said  it  is  evident  that  a 
monobasic  acid  can  form  only  normal  salts,  with  monoacid 
bases ;  whereas  with  polyacid  bases  it  may  form  basic  salts  as 
well  as  normal  salts.  Again,  polybasic  acids  may  form  either 
normal  salts,  acid  salts,  or  basic  salts  with  polyacid  bases; 
whereas  with  monoacid  bases  they  can  form  only  acid  salts  and 
normal  salts. 

Acidimetry  and  Alkalimetry.  —  The  fact  that  it  always  takes 
a  definite  amount  of  acid  and  a  definite  amount  of  base  to 
exactly  neutralize  each  other  is  used  in  the  quantitative  esti- 
mation of  acids  and  markedly  alkaline  bases  in  processes  that 
are  called  acidimetry  and  alkalimetry.  A  given  volume  of  a 
solution  of  an  acid  of  known  strength  will  always  neutralize 
a  perfectly  definite  volume  of  a  given  solution  of  an  alkali.  If 
we  place  the  acid  of  known  strength  in  the  burette  A  (Fig.  33) 
and  an  alkaline  solution  of  unknown  strength,  e.g.  of  sodium 
hydroxide,  in  the  burette  B,  then  run  out  say  20  cc.  into  the 
dish  D  and  find  that  it  is  necessary  to  add  23.6  cc.  of  the  acid 
in  A  to  just  make  the  solution  in  the  dish  D  neutral  to  litmus, 


ACIDS,    BASES,    SALTS,    CHEMICAL   EQUILIBRIUM 


129 


it  is  possible  to  compute  the  strength  of  the  sodium  hydroxide 
solution  from  the  data  at  hand.  Any  solution  of  known  strength 
is  called  a  standard  solution.  In  working  with  acid  solutions 
of  known  strength  normal 
solutions  are  frequently  used. 
A  normal  solution  of  an  acid  is 
one  which  contains  1.008  grams 
of  replaceable  hydrogen  per  liter 
of  solution.  Thus  a  normal 
solution  of  hydrochloric  acid 
would  contain  36.468  grams  of 
pure  HC1  per  liter,  for  in  this 
amount  there  are  1.008  grams 

of  replaceable  hydrogen.  In  ^j^jtz  '  S^3  A"- 
the  case  of  hydrobromic  acid,  a 
liter  of  normal  solution  would 
contain  80.928  grams  of  HBr;  |"j— B 
in  the  case  of  sulphuric  acid  a 
liter  of  normal  solution  would 
contain  49.043  grams  of  H2SO4, 
etc.  Sometimes  solutions  of 
one  half,  one  tenth,  or  one 
twentieth,  etc.,  of  the  strength 
of  normal  solutions  are  em- 
ployed; these  contain  the  cor- 
responding amounts  of  re- 


FIG.  33. 


placeable  hydrogen  per   liter. 

Solutions  of   an  acid  may  be 

prepared  which  contain  some 

multiple    of    1.008    grams    of 

replaceable  hydrogen  or  some 

aliquot   part  thereof   per  liter.      These  are  then  called  twice 

normal,  normal,  half  normal,  fiftieth  normal,  etc.,  as  the  case 

may  be. 

A  normal  solution  of  an  alkali  is  one  that  will  just  neutralize  a 
normal  solution  of  an  acid  volume  for  volume ;  consequently  a 
normal  solution  of  an  alkali  is  a  solution  that  contains  the  chemical 
equivalent  of  1.008  grams  of  replaceable  hydrogen  per  liter.  Al- 
kaline solutions  may  be  made  up  as  normal  solutions  or  as 
some  multiple  or  fractional  part  of  normal,  just  as  in  the  case 


130  OUTLINES   OF   CHEMISTRY 

of  the  acid  solutions.  So,  for  instance,  a  normal  solution  of 
sodium  hydroxide  NaOH  contains  40.008  grams  of  pure  NaOH 
per  liter,  and  a  tenth  normal  solution  contains  one  tenth  as 
much  per  liter. 

•The  process  of  exactly  neutralizing  an  acid  solution  with  an 
alkaline  solution  to  ascertain  the  strength  of  one  of  them  in 
terms  of  the  other,  as  illustrated  in  connection  with  Fig.  33,  is 
called  titration.  If,  in  the  instance  cited  above,  the  acid  solu- 
tion was  normal,  then  the  20  cc.  of  NaOH  solution  required 
would  just  be  equal  to  23.6  cc.  of  normal  solution.  In  other 

23  6 
words,  the  NaOH  solution  would  be  — '—  normal,  and  contain 

23  6 

-— ^-  x  40.00  grams  of  pure  NaOH,  or  47.20  grams  per  liter, 
.zu 

Indicators.  — Litmus  may  serve  to  indicate  the  acidity  or 
alkalinity  of  a  solution.  There  are,  however,  many  other 
coloring  matters  that  change  their  hue  on  being  treated  with 
an  acid  and  an  alkali  in  succession.  Such  substances  may  con- 
sequently also  serve  as  indicators.  So  phenolphthalein  is  color- 
less in  acid  solution,  pink  in  a  faintly  alkaline  solution,  and  a 
beautiful  purplish  red  in  strongly  alkaline  solution.  Methyl 
orange  is  red  when  acid,  straw  yellow  when  alkaline,  and 
orange  colored  w^hen  neutral.  Congo  red  is  red  when  alkaline 
and  blue  when  acid ;  that  is,  it  acts  in  just  the  opposite  way  that 
litmus  does.  Turmeric  paper,  that  is  paper  soaked  in  a  decoc- 
tion of  turmeric  root,  is  yellow  when  neutral  or  acid,  but  turns 
brown  when  moistened  with  an  alkaline  solution.  There  are 
still  other  indicators  in  use,  but  those  mentioned  are  the  ones 
commonly  employed  in  the  laboratory. 

Hydrolysis.  —  When  phosphorus  trichloride  is  brought  into 
contact  with  water,  violent  action  ensues,  heat  is  liberated,  and 
phosphorous  and  hydrochloric  acids  are  formed,  thus :  — 

PC13  +  3  H20  =  P(OH)3  +  3  HC1. 

The  action  is  complete  and  irreversible.  No  such  change 
takes  place,  for  example,  when  sodium  chloride  NaCl  is  treated 
with  water.  The  solution  in  this  case  remains  quite  neutral, 
and  the  salt  may  be  recovered  by  evaporating  off  the  water.  In 
phosphorus  trichloride  we  have  a  compound  which  in  contact 
with  water  readily  passes  over  into  the  much  more  stable  com- 


ACIDS,    BASES,    SALTS,    CHEMICAL   EQUILIBRIUM  131 

pounds  hydrochloric  and  phosphorous  acids.  In  the  forma- 
tion of  hydrochloric  acid,  the  great  affinity  of  hydrogen  for 
chlorine  comes  into  play,  and  in  the  formation  of  phosphorous 
acid  the  strong  affinity  of  phosphorus  for  oxygen  exerts  itself. 
We  thus  see  that  all  three  chlorine  atoms  in  PC13  are  exchanged 
for  OH  groups  taken  from  three  water  molecules,  the  remain- 
ing hydrogen  atoms  of  which  unite  with  the  chlorine  atoms  to 
form  hydrochloric  acid.  When  a  substance  is  thus  decomposed  by 
water,  the  process  is  termed  hydrolytic  decomposition,  or  hydrolysis. 
Very  many  salts  suffer  hydrolysis  in  water  to  a  slight  extent, 
others  are  not  decomposed  by  water,  and  still  others  are  com- 
pletely hydrolyzed.  Thus  we  have  seen  that  PC13,  which  is  a 
salt  of  an  extremely  weak  basic  element,  phosphorus,  with  a 
very  strong  acidic  element,  chlorine,  is  completely  decomposed  by 
hydrolysis.  On  the  other  hand,  sodium  chloride,  a  salt  of  the 
strongly  basic  sodium  with  the  strongly  acidic  chlorine,  is  not 
hydrolyzed.  In  general,  salts  of  weak  bases  with  strong  acids 
are  more  or  less  hydrolyzed  when  brought  into  contact  with  water. 
The  same  is  true  of  salts  of  strong  bases  with  very  weak  acids.  So 
cupric  sulphate  CuSO4,  ferric  chloride  FeCl3,  and  in  general  all 
the  ordinary  salts  of  the  heavy  metals,  are  somewhat  hydrolytically 
decomposed  when  dissolved  in  water.  This  is  made  evident,  for 
instance,  by  the  fact  that  all  these  solutions  react  acid  toward 
litmus/  On  the  other  hand,  salts  like  sodium  carbonate  Na2CO8 
and  borax  Na2B4O7,  which  contain  a  very  weak  acid  radical 
united  to  the  strongly  basic  sodium,  are  also  hydrolyzed  to  some 
extent.  This  is  evidenced  by  the  fact  that  their  solutions  react 
alkaline  toward  litmus.  Even  sodium  bicarbonate  NaHCO3  is 
thus  hydrolyzed,  and  its  solutions  react  alkaline  toward  litmus,' 
though  the  salt  still  contains  hydrogen  that  is  replaceable  by  a 
metal  and  consequently  is  an  acid  salt.  Thus  it  is  that  the 
normal  salts  need  not  necessarily  yield  solutions  that  are  neutral 
toward  indicators,  for  many  of  them  are  hydrolyzed.  Indeed, 
neutrality  toward  indicators  is  met  only  in  dealing  with  solutions 
of  normal  salts  of  strong  acids  with  strong  bases.  Thus,  the 
chlorides,  sulphates,  nitrates,  and  bromides  of  sodium,  potassium, 
calcium,  and  magnesium  in  solution  are  neutral  toward  indica- 
tors ;  whereas  the  corresponding  salts  of  iron,  copper,  lead,  and 
mercury  have  acid  reactions  in  solution,  and  the  carbonates, 
silicates,  and  borates  of  sodium  and  potassium  yield  alkaline  so- 


182  OUTLINES   OF   CHEMISTRY 

lutions.  From  this  it  is  evident  that  the  processes  of  acidimetry 
and  alkalimetry  described  above  can  be  used  only  in  estimat- 
ing the  strength  of  solutions  of  fairly  pronounced  acids  and 
alkalies. 

If  now  we  take  cupric  chloride  CuCl3  and  dissolve  it  in  water, 
we  find  that  the  solution  is  distinctly  acid  toward  litmus  and 
other  indicators,  showing  that  the  salt  has  suffered  decomposi- 
tion to  a  certain  degree.  The  hydrolysis  which  has  taken  place 
has  resulted  in  the  liberation  of  a  slight  amount  of  hydrochloric 
acid.  One  may  indicate  the  change,  at  least  in  part,  thus :  — 

CuCl2  +  H2O^±Cu(OH)Cl  +  HC1 ; 

meaning  that  only  a  small  percentage  of  the  salt  is  thus  hydro- 
lyzed  in  solutions  that  contain  say  ten  per  cent  or  more  of  the 
salt. 

Such  solutions  remain  clear,  the  basic  cupric  chloride  forming 
with  the  cupric  chloride  CuCl2  a  compound  which  consists  of  a 
soluble,  though  slightly  basic  cupric  chloride.  On  dilution  of 
the  solution  of  cupric  chloride,  the  hydrolysis  proceeds  further 
and  a  basic  cupric  chloride  finally  forms,  which  is  richer  in  base 
and  poorer  in  chlorine  than  that  in  the  stronger  solutions.  This 
basic  salt  is  difficultly  soluble  and  gradually  separates  out  in 
form  of  a  precipitate.  The  more  dilute  the  solution,  then,  the 
more  does  the  reaction  proceed  in  the  direction  indicated  by 
the  upper  arrow  in  the  equation.  On  the  other  hand,  on  con- 
centrating the  solution  the  reaction  proceeds  from  right  to 
left;  i.e.  it  reverses.  We  have  here  then  a  case  of  reversible 
hydrolysis. 

Mass  Action ;  Chemical  Equilibrium.  —  Thus  we  see,  in  the 
instance  just  mentioned,  that  the  more  water  there  is  added  to 
the  cupric  chloride  solution,  the  greater  is  the  extent  of  the 
hydrolysis.  The  water  consequently  influences  the  process  to 
proceed  from  left  to  right  (see  the  equation) ;  but  it  is  to  be 
borne  in  mind  that  it  is  the  mass  of  water  present  relative  to 
the  amount  of  cupric  chloride  in  the  solution  that  really  deter- 
mines the  direction  of  the  reaction. 

The  amount  of  matter  contained  in  unit  of  volume  is  commonly 
called  the  concentration.  In  the  case  of  a  solution,  the  concen- 
tration is  the  amount  of  dissolved  substance  contained  in  unit 
of  volume  of  the  solution.  Very  commonly  the  concentration 


ACIDS,    BASES,    SALTS,    CHEMICAL  EQUILIBRIUM  133 

of  a  solution  is  expressed  by  stating  the  number  of  gram- 
molecules  contained  in  one  liter  of  the  solution.  By  a 
gram-molecule  is  meant  the  molecular  weight  in  grams.  Thus 
a  gram-molecule  of  hydrochloric  acid  is  36.46  grams  HC1;  a 
gram-molecule  of  potassium  iodide,  166.02  grams  KI. 

We  may  state  the  facts  elucidated  in  the  case  of  the  hydroly- 
sis of  cupric  chloride  by  saying  that  the  extent  of  the  hy- 
drolysis is  determined  by  the  relative  concentrations  of  the 
cupric  chloride  and  the  water  in  the  solution ;  i.e.  by  the  relative 
masses  of  the  substances  that  are  acting  on  each  other.  The 
amount  of  undecomposed  cupric  chloride,  water,  basic  cupric 
chloride,  and  hydrochloric  acid  which  the  solution  contains  at 
any  particular  temperature  and  concentration  of  the  cupric 
chloride  solution  is  perfectly  definite  ;  and  the  amounts  of  these 
four  ingredients  are  said  to  be  in  chemical  equilibrium  with  one 
another.  The  equilibrium  of  this  system  may  be  disturbed  by 
changing  the  concentration  of  any  one  of  the  four  ingredients 
that  make  up  the  solution.  Thus  to  abstract  water  from  the 
solution  causes  the  action  to  proceed  in  the .  direction  of  the 
lower  arrow,  — 

CuCl2  +  H20  5>  Cu(OH)Cl  +  HC1 ; 

while  to  add  water  causes  the  action  to  go  in  the  direction 
of  the  upper  arrow.  To  abstract  cupric  chloride  from  the 
system  causes  the  equilibrium  to  be  displaced  in  the  direction 
of  the  lower  arrow,  for  this  practically  amounts  to  the  same 
thing  as  adding  more  water  relatively.  Addition  of  cupric 
chloride  causes  the  opposite  effect.  Addition  of  hydrochloric 
acid  to  the  system  causes  the  action  to  proceed  in  the  direction 
of  the  lower  arrow  ;  abstracting  hydrochloric  acid  causes  the 
equilibrium  to  be  displaced  in  the  direction  of  the  upper  arrow. 
Addition  of  basic  cupric  chloride  to  the  system  causes  the  equi- 
librium to  be  changed  in  the  direction  of  the  lower  arrow,  while 
removal  of  basic  cupric  chloride  from  the  system  causes  the 
reaction  to  proceed  in  the  direction  of  the  upper  arrow.  It  is 
obvious  that  if  either  the  hydrochloric  acid  or  the  basic  cupric 
chloride  were  taken  from  the  system  as  fast  as  formed,  the  re- 
action would  go  to  completeness  from  left  to  right  and  with 
increased  rapidity. 

What  has  been  thus  presented  is  really  a  special  case  of  a 


134  OUTLINES   OF   CHEMISTRY 

general  law,  termed  the  law  of  mass  action,  which  may  be  stated 
thus  :  — 

The  speed  or  rate  of  any  chemical  change  is  proportional  to  the 
active  mass,  that  is,  the  molecular  concentration  of  each  substance 
engaged  in  the  reaction.  This  is  universal  and  holds  for  all 
chemical  changes,  whether  they  are  reversible  or  not.  In 
case  of  reversible  reactions,  the  law  holds  for  the  change 
from  right  to  left  as  well  as  from  left  to  right,  and  hence  the 
final  chemical  equilibrium  reached  is  also  determined  by  the  law 
of  mass  action.  One  can  best  comprehend  this  by  thinking  of 
the  equilibrium  as  reached  when  the  rate  of  speed  of  the  for- 
ward action  just  equals  that  of  the  reverse  action.  Chemical 
equilibrium  is  commonly  regarded  as  dynamic  rather  than  static 
in  character. 

Additional  Illustrations  of  Chemical  Equilibrium  and  the  Oper- 
ation of  the  Law  of  Mass  Action.  —  In  the  first  chapter  it  was 
stated  that  the  factors  which  determine  whether  a  chemical 
change  will  go  on  or  not  are  :  (1)  the  right  substances  must  be 
brought  into  contact,  i.e.  chemical  attraction  or  chemical  affinity 
must  exist  between  the  substances  that  are  to  react ;  (2)  the 
temperature  must  be  properly  chosen  ;  (3)  the  pressure  is  of 
consequence,  particularly  when  a  gas  enters  into  the  change  ; 
(4)  the  concentrations  of  the  active  substances  must  be  con- 
sidered. All  of  these  factors  determine  not  only  whether  the 
change  will  proceed  at  all  or  not,  but  they  also  influence  the 
rate  with  which  the  action  proceeds  and  consequently  affect 
the  final  equilibrium  reached.  Now  it  is  clear  that  it  is  with 
factor  (4),  above  mentioned,  that  the  law  of  mass  action  is 
concerned. 

There  are  many  reactions  which  are,  so  far  as  we  know, 
irreversible  ;  that  is,  they  go  to  completion  in  one  direction. 
We  have  already  seen  that  the  hydrolysis  of  phosphorus  chlo- 
ride is  of  this  class.  The  combustion  of  calcium,  magnesium, 
or  sodium  in  oxygen,  the  decomposition  of  potassium  chlorate 
into  potassium  chloride  and  oxygen,  the  neutralization  of  po- 
tassium hydroxide  by  hydrochloric  acid,  the  burning  of  sugar 
to  water  and  carbon  dioxide,  are  further  examples  of  this  kind. 
In  these  reactions,  the  chemical  affinity  factor,  namely  (1)  above, 
is  really  the  determining  one  ;  i.e.  its  influence  overshadows  all 
the  other  factors  very  greatly,  and  so  the  action  goes  on  in  one 


ACIDS,   BASES,    SALTS,   CHEMICAL  EQUILIBRIUM  135 

direction  to  completion  and  is  irreversible.  The  cases  of  irre 
versible  reactions  are  after  all  then  fairly  clearly  distinguished, 
for  as  a  rule  they  belong  to  one  of  two  categories  ;  namely : 
(1)  they  represent  the  formation  of  very  stable  compounds 
directly  from  the  elements,  in  which  processes  powerful  affinities 
come  into  play  ;  or  (2)  they  represent  the  decomposition  of 
relatively  unstable  or  complicated  compounds  into  simpler  and 
much  stabler  ones. 

The  burning  of  barium  to  barium  oxide  is  a  typical  illustra- 
tion of  the  first  class;  the  decomposition  of  sugar  or  nitroglycer- 
ine by  heat  illustrates  the  second  class.  In  speaking  of  irre- 
versible reactions,  it  must  be  borne  in  mind  that  the  term  does 
not  mean  that  the  original  substances  taken  cannot  be  got 
back  by  roundabout  means.  So  while  the  burning  of  magne- 
sium to  magnesium  oxide  is  an  irreversible  reaction,  it  is  yet 
possible  to  get  back  the  metallic  magnesium  and  also  the  oxygen 
that  it  contains.  In  this  sense,  of  course,  all  chemical  reactions 
could  be  reversed,  for  matter  cannot  be  destroyed,  but  simply 
transformed.  What  we  mean  by  an  irreversible  reaction,  in  the 
sense  in  which  the  term  is  here  used,  is  a  reaction  that  cannot  be  re- 
versed entirely  or  in  part  by  simply  altering  the  temperature, 
pressure,  or  the  concentration  of  the  substances  concerned  in  the 
reaction. 

In  the  irreversible  reactions  the  factors  of  temperature, 
pressure,  and  concentration  cannot  reverse  the  process.  But 
in  many  chemical  changes  the  affinities  that  come  into  play  in  fix- 
ing the  direction  in  which  the  change  will  go  on  are  so  well  balanced 
that  changes  of  temperature,  pressure,  or  concentration  suffice  as 
determining  factors  in  altering  the  direction  the  reaction  takes. 
Such  reactions  are  consequently  reversible.  This  class  of  reactions 
is  very  large  indeed.  It  is  therefore  evident  that  at  constant 
temperature  and  pressure  the  effect  of  concentration  is  of  vast 
importance  in  case  of  reversible  reactions,  for  in  these  it  de- 
termines the  direction  of  the  change  and  consequently  the 
final  equilibrium.  On  the  other  hand,  in  the  irreversible  re- 
actions the  concentration  changes  can  only  affect  the  rate  of 
the  reaction.  Thus,  in  the  burning  of  magnesium  ribbon  the 
final  product  is  MgO,  and  whether  the  action  proceeds  in 
oxygen  at  atmospheric  pressure,  or  in  compressed  oxygen,  only 
affects  the  rate  of  the  combustion,  not  the  character  or  the 


136  OUTLINES   OF  CHEMISTRY 

amount  of  the  final  product.  But  when,  for  instance,  chlorine 
acts  on  water  in  diffused  light,  we  have  a  case  of  a  reversible 
reaction,  thus:  — 

H2O  +  C12^±HOC1  +  HC1. 

All  four  ingredients  are  finally  in  equilibrium  at  any  given 
temperature  and  pressure.  By  increasing  either  the  relative 
amount  of  water  or  chlorine  the  change  progresses  somewhat 
more  from  left  to  right ;  the  reverse  happens  by  increasing  the 
relative  concentration  of  either  the  hydrochloric  acid  or  hypo- 
chlorous  acid.  By  diminishing  the  concentration  of  either  the 
water  or  chlorine  or  both,  the  reaction  proceeds  from  right  to 
left.  Decreasing  the  concentration  of  either  the  hypochlorous 
acid  or  hydrochloric  acid  or  both  causes  the  change  to  proceed 
from  left  to  right.  If  we  were  to  abstract  say  the  hypochlo- 
rous acid  as  fast  as  it  forms,  the  reaction  would  complete  itself 
from  left  to  right.  Now,  in  the  sunlight  hypochlorous  acid 
undergoes  decomposition,  thus:  — 

2  HOC1  =  O2  +  2  HC1. 

Therefore  as  the  oxygen  escapes  and  only  hydrochloric  acid 
remains  in  the  solution,  we  have  (when  chlorine  acts  on  water 
in  sunlight)  a  reaction  which  goes  on  to  completion.  This 
reaction  is  complete  because  of  the  removal  of  one  of  the  in- 
gredients ;  namely,  the  hypochlorous  acid. 

Again,  when  sulphuric  acid  acts  on  common  salt  in  moder- 
ately dilute  solution,  10  to  20  per  cent  for  instance,  an  equi- 
librium is  established  which  may  be  expressed  thus  :  — 

NaCl  +  H2S04^±NaHSO4  +  HCL 

The  action  is  reversible,  for  it  is  possible  to  displace  the  equilib- 
rium in  the  one  direction  or  the  other  by  changing  the  concen- 
tration of  the  substances  that  enter  into  the  change.  Now, 
when  concentrated  sulphuric  acid  is  poured  on  sodium  chloride 
and  the  mass  becomes  warm,  the  reaction  will  complete  itself 
from  left  to  right;  for  the  hydrochloric  acid  is  volatile,  and 
under  the  conditions  of  the  experiment  it  can  escape  and  so 
get  out  of  the  field  of  action.  This  does  not  necessarily  mean 
that  the  sulphuric  acid  is  stronger  than  the  hydrochloric  acid 
and  so  drives  the  latter  out,  as  was  formerly  supposed.  It  will 


ACIDS,   BASES,    SALTS,   CHEMICAL   EQUILIBRIUM  137 

be  observed  that  the  determining  factor  is  rather  the  volatility 
of  the  hydrochloric  acid,  which  takes  it  out  of  the  reacting 
mass.  Indeed,  it  is  possible  to  displace  the  hydrochloric  acid 
from  common  salt  by  boiling  it  with  strong  solutions  of  much 
weaker  acids  than  hydrochloric  acid,  provided  that  the  acid  so 
employed  is  non-volatile.  So,  for  example,  it  is  possible  to 
evolve  hydrochloric  acid  from  salt  by  employing  boric  acid, 
which,  as  we  shall  learn,  is  a  very  weak  yet  practically  non-vola- 
tile acid. 

Whenever  liquids  act  on  liquids,  or  solids  act  on  liquids, 
forming  a  product  which  is  gaseous  and  so  escapes,  the  reaction 
proceeds  practically  to  completion.  The  same  is  true  whenever 
in  such  cases  a  solid  forms  which  is  insoluble,  i.e.  is  practically 
not  acted  upon,  and  so  is  thrown  out  of  the  field  of  action. 
Thus,  for  instance,  when  sodium  sulphate  acts  on  barium  chloride 
we  have  the  following  change  taking  place  :  — 

BaCl2  +  Na2S04  =  BaSO4  +  2  NaCL 

This  goes  practically  to  completion  because  the  barium  sulphate 
formed  is  very  difficultly  soluble,  and  nearly  all  of  it  drops  out 
of  the  field  of  action  as  a  precipitate. 

In  the  case  of  gases  we  frequently  have  instances  of  rever- 
sible changes.  At  red  heat,  water  vapor  partially  decomposes 
into  hydrogen  and  oxygen;  at  still  higher  temperatures,  the 
reaction  progresses  further  in  the  sense  mentioned,  whereas  on 
cooling  it  again  is  reversed.  The  process  of  thus  decomposing 
a  substance  on  heating  it  is  called  dissociation.  It  was  studied 
particularly  by  Henri  Saint  Claire  Deville.  We  shall  consider 
cases  of  the  dissociation  of  gases  more  carefully  later. 

Strength  of  Acids  and  Bases.  — The  relative  strengths  of  acids 
has  been  a  favorite  subject  of  study  with  chemists.  By  having, 
let  us  say,  tenth  normal  solutions  of  hydrochloric,  sulphuric, 
and  acetic  acids  each  separately  act  on  a  piece  of  zinc  (the 
pieces  being  arranged  so  as  to  expose  the  same  area  of  zinc  to 
each  acid)  and  estimating  the  volume  of  hydrogen  liberated  by 
each  acid  per  minute,  it  is  possible  to  compare  the  relative 
strengths  of  the  acids.  The  apparatus  for  this  purpose  might 
be  arranged  as  in  Fig.  14.  In  each  tube  is  placed  a  piece  of 
zinc  of  the  same  size  and  shape.  The  whole  apparatus  is  then 
filled  with  water,  the  same  quantity  being  used  in  each  case. 


138  OUTLINES   OF  CHEMISTRY 

The  acids  are  then  introduced  in  chemically  equivalent  amounts 
from  above  by  means  of  the  stopcocks,  care  being  taken  to 
admit  no  air  through  the  cocks.  The  volumes  of  hydrogen 
evolved  per  minute  may  then  readily  be  read.  We  should 
thus  be  estimating  the  strengths  of  these  acids  by  their  rate  of 
action  upon  zinc.  It  is,  of  course,  possible  to  use  other 
characteristic  activities  of  acids  as  a  basis  of  estimating  their 
strength.  Similarly,  strengths  of  alkalies  might  be  compared 
by  measuring  the  rate  with  which  they  transform  a  fat  into 
soap  (which  see). 


CHAPTER   X 

NITROGEN,    THE   ATMOSPHERE,    AND  THE   ELEMENTS 
OP   THE   HELIUM   GROUP 

History  and  Occurrence  of  Nitrogen.  —  In  1772  Dr.  Rutherford, 
professor  of  botany  at  Edinburgh,  found  that  when  animals  are 
confined  in  an  air-tight  space,  the  air  they  breathe  becomes 
incapable  of  supporting  combustion  or  respiration.  After  treat- 
ing such  air  with  caustic  potash  solution  to  absorb  the  carbon 
dioxide,  then  called  "  fixed  air,"  he  showed  that  the  remaining 
gas  supported  neither  life  nor  combustion.  .  A  lighted  candle 
thrust  into  the  gas,  for  instance,  was  immediately  extinguished. 
He  called  this  residual  gas  "  mephitic  air."  Priestley  burned 
carbon  in  a  confined  volume  of  air,  and  then  treated  the  latter 
with  limewater ;  thus  the  carbon  dioxide  formed  during  the  com- 
bustion was  absorbed,  and  a  residual  gas  was  obtained,  which  he 
called  " phlogisticated  air."-  He  found  that  one  fifth  of  the  vol- 
ume of  atmospheric  air  can  thus  be  converted  into  "  fixed  air  " 
and  absorbed  by  caustic  lime.  But  he  did  not  regard  the  "  phlo- 
gisticated "  air  he  had  prepared  as  a  constituent  of  the  atmos- 
phere. It  was  Scheele  (1777)  who  first  showed  that  there  are 
two  different  gases  in  the  air.  Lavoisier  was  the  first  to  con- 
sider mephitic  or  phlogisticated  air  as  an  element.  He  called  it 
azote,  because  of  its  inability  to  support  life.  The  name  nitro- 
gen was  given  to  the  gas  by  Chaptal,  because  it  forms  an  essen- 
tial constituent  of  niter  or  saltpeter.  Cavendish  showed  that 
nitrogen  obtained  from  air  is  essentially  a  simple  body  which  is 
somewhat  lighter  than  ordinary  air ;  and,  indeed,  till  1894  the 
residual  gas  thus  prepared  was  regarded  as  pure  nitrogen.  Sir 
William  Ramsay  and  Lord  Rayleigh  showed  that  the  gas  remain- 
ing after  the  oxygen  and  carbon  dioxide  have  been  removed 
from  the  air  consists  of  98.814  per  cent  nitrogen  and  1.186  per 
cent  other  gases,  which,  unlike  nitrogen,  will  not  unite  with 
oxygen  or  with  red-hot  magnesium.  This  notable  observation 
led  to  the  discovery  of  the  new  elements  of  the  helium  group. 

139 


140 


OUTLINES   OF   CHEMISTRY 


About  80  per  cent  of  the  volume  of  atmospheric  air  consists 
of  nitrogen  in  the  free  state.  In  combination  with  carbon, 
hydrogen,  and  oxygen,  nitrogen  forms  an  essential  constituent 
of  the  bodies  of  all  plants  and  animals.  It  is  found  especially 
in  the  blood,  muscles,  nerves,  seeds,  and,  in  general,  in  all  tissues 
that  are  concerned  in  movement  or  reproduction.  When  plants 
and  animals  die  and  their  bodies  decay,  their  nitrogen  content 
passes  over  into  simpler  compounds,  namely,  ammonia,  nitrites, 
and  nitrates  (which  see).  Thus  it  is  that  in  all  soils  nitrogen 
is  present  in  the  form  of  nitrates  and  ammonium  salts.  It  also 
occurs  in  all  refuse  matter  of  plant  or  animal  origin,  like  barn- 
yard manure,  guano,  sewage,  etc.  In  coal,  which  represents 
the  remains  of  plants  of  the  carboniferous  age,  nitrogen  is  found 
in  combination  with  hydrogen,  carbon,  and  oxygen.  In  minute 
quantities,  nitrogen  also  occurs  in  granitic  rocks,  in  meteoric 
iron,  and  in  steel.  In  Chili  saltpeter,  consisting  chiefly  of 
sodium  nitrate,  nitrogen  occurs  in  large  quantities. 

Preparation  and  Properties  of  Nitrogen.  —  Nitrogen  which  is 
approximately  99  per  cent  pure  may  be  prepared  from  the  air 

by  removing  the  oxygen  from  the 
latter.  This  is  generally  accom- 
plished by  heating  in  the  air  some 
elementary  substance  which  will  read- 
ily combine  with  oxygen,  forming  an 
oxide  that  is  either  a  non-volatile 
solid  or  that  can  readily  be  removed 
by  absorption  in  some  liquid.  Thus, 
when  phosphorus  is  burned  in  a  little 
dish  resting  on  water  under  a  bell 
jar  (Fig.  34),  phosphorus  pentoxide 

is  formed,  which  is  a  solid  that  is  readily  absorbed  by  water, 
forming  phosphoric  acid  :  — 

P4  +  5  O2  =  2  P2O6,  and 
3  H20  +  P205  =  2  H3P04. 

Again,  air  may  be  passed  over  red-hot  copper,  when  the  latter 
unites  with  the  oxygen,  forming  cupric  oxide  CuO,  which  is 
non- volatile,  thus  leaving  the  nitrogen.  The  oxygen  may  also 
be  removed  from  the  air  by  shaking  the  latter  with  an  alkaline 
solution  of  pyrogallic  acid,  which  readily  absorbs  oxygen,  and 


NITROGEN,   AIR,   AND  THE   HELIUM   GROUP  141 

which  is  frequently  used  for  this  reason  in  gas  analysis.  Left 
in  contact  with  moist  yellow  phosphorus,  the  air  is  also  deprived 
of  its  oxygen  even  at  room  temperatures.  This  fact  is  often 
used  in  determining  the  amount  of  oxygen  in  a  given  sample 
of  gas.  By  cooling  air  to  —  182°  the  oxygen  liquefies,  leaving 
the  nitrogen  in  form  of  a  gas. 

Pure  nitrogen  cannot  very  well  be  prepared  from  atmospheric 
air,  for  the  gases  of  the  helium  group,  with  which  it  is  always 
contaminated,  are  chemically  very  inert,  and  hence  difficult  to 
remove.  Pure  nitrogen  is  prepared  from  compounds  in  which 
it  occurs.  Thus,  by  treating  ammonia  NH3  with  chlorine, 
nitrogen  and  hydrochloric  acid  are  formed,  the  latter  uniting 
with  some  of  the  ammonia  (which  should  be  present  in  excess) 
to  form  ammonium  chloride,  which  dissolves  in  water.  The 
reactions  may  be  expressed  thus :  — 

2  NH3  +  3  C12  =  6  HC1  +  N2. 
NH3  +  HC1  =  NH4C1. 

The  simplest  way  of  preparing  pure  nitrogen  consists  of  heating 
ammonium  nitrite  NH4NO2,  either  in  pure  form  or  in  strong 
aqueous  solution.  The  compound  when  thus  treated  decom- 
poses into  water  and  nitrogen  :  — 

NH4N02=2H20  +  N2. 

Frequently  ammonium  nitrite  is  not  at  hand,  and  a  mixture  of 
sodium  nitrite  and  either  ammonium  chloride  or  ammonium 
sulphate  is  employed.  By  the  interaction  of  the  sodium  nitrite 
and  the  ammonium  salt  employed,  ammonium  nitrite  is  formed, 
which  on  heating  decomposes  into  water  and  nitrogen.  When, 
for  instance,  sodium  nitrite  and  ammonium  sulphate  are  em- 
ployed, the  reaction  is  as  follows  :  — 

2  NaN02  +  (NH4)2S04  =  Na2SO4  +  4  H2O  +  2  N2. 

By  heating  ammonium  bichromate  (NH4)2Cr2O7,  or  a  mixture 
of  ammonium  chloride  and  potassium  bichromate,  nitrogen  is 
formed,  thus :  — 

K2Cr207  +  2  NH4C1  =  2  KC1  +  (NH4)2Cr2O7, 
and  (NH4)2Cr207  =  CraO,  +  4  H2O  +  N2 ; 

or,  by  combining  the  two  equations, — 

K2Cr207  +  2  NH4C1  =  2  KC1  +  CraO8  +  4  H2O  +  N2. 


142  OUTLINES   OF   CHEMISTRY 

When  oxides  of  nitrogen  are  passed  over  red-hot  copper,  cupric 
oxide  and  nitrogen  are  formed,  for  example  :  — 


When  urea  CO(NH2)2  is  oxidized  by  means  of  hypochlorous 
or  hypobromous  acids  or  their  salts,  nitrogen  is  formed.  So, 
for  instance,  with  potassium  hypobromite  the  reaction  is  :  — 

CO(NH2)2  +  3  KBrO  =  2  H2O  +  3  KBr  +  CO2  +  N2. 

The  potassium  hypobromite  solution  as  usually  prepared  con- 
tains an  excess  of  caustic  potash,  which  at  once  absorbs  the 
carbon  dioxide,  forming  potassium  carbonate,  which  dissolves 
in  water:  — 

2  KOH  +  C02  =  K2C03  +  H20. 

In  estimating  the  quantity  of  urea  in  urine,  which  often  needs 
to  be  done  in  medical  practice,  these  reactions  are  used. 

Nitrogen  is  a  colorless,  odorless,  tasteless  gas,  which  is  0.9672 
time  as  heavy  as  air.  It  may  be  liquefied  and  solidified.  Liquid 
nitrogen  is  colorless,  and  boils  at  —  195.5°  at  atmospheric  pres- 
sure. The  critical  temperature  is  —  146°,  at  which  it  requires 
a  pressure  of  35  atmospheres  to  liquefy  the  gas.  Liquid  nitro- 
gen has  the  specific  gravity  0.80  at  its  boiling  point.  Solid 
nitrogen  is  a  white,  crystalline  substance  melting  at  —  214°; 
its  specific  gravity  is  1.0265  at  —  252.5°.  Nitrogen  is  less 
soluble  in  water  than  oxygen.  At  10°,  1000  cc.  of  water  dis- 
solve 16.1  cc.  of  nitrogen,  while  at  0°,  20.34  cc.  are  absorbed. 

At  ordinary  temperatures,  nitrogen  is  a  very  inert  element 
chemically.  At  higher  temperatures  it  unites  with  '  lithium, 
boron,  silicon,  magnesium,  barium,  strontium,  or  calcium  to  form 
nitrides.  Lithium  burns  readily  in  nitrogen,  and  even  unites 
slowly  with  that  gas  at  ordinary  temperatures,  forming  the 
nitride  Li3N.  Magnesium  at  red  heat  absorbs  nitrogen  greedily, 
forming  Mg8N2.  In  general,  nitrogen  is  trivalent  in  the  nitrides. 
When  nitrogen  and  oxygen  are  mixed  and  subjected  to  the 
action  of  the  electric  spark  (Fig.  35),  nitrogen  and  oxygen 
unite  to  form  an  oxide  of  a  brown  color.  Its  formula  is  NO2  ; 
at  room  temperatures  it  is  N2O4.  Hydrogen  and  nitrogen  when 
mixed  and  similarly  sparked  yield  small  amounts  of  ammonia, 
which  is  a  nitride  of  hydrogen  having  the  composition  NH3. 


NITROGEN,   AIR,   AND   THE   HELIUM   GROUP 


143 


FIG.  35. 


Due  to  electrical  disturbances  in  the  atmosphere,  especially 
during  thunder  storms  when  lightning  flashes  from  cloud  to 
cloud  or  to  earth,  small  amounts  of  ammonia  and  oxides  of 
nitrogen  are  formed. 

The  atomic  weight  of  nitrogen  is  14.01 ;  and  since  at  0°  and 
760  mm.  pressure  22.38  liters  of  nitrogen  weigh  27.98  grams, 
the  molecule  contains 
2  atoms  and  the  mo- 
lecular formula  is  N2. 
This  is  also  in  harmony 
with  the  composition 
of  ammonia  and  of  the 
oxides  of  nitrogen  by 
volume,  as  will  appear 
later.  In  compounds 
nitrogen  is  either  triv- 
alent  or  pentavalent. 
Its  atomic  weight  was 
determined  by  Stas, 
who  ascertained  the 
proportion  by  weight  in  which  nitrogen  exists  in  silver  nitrate 
and  in  ammonium  chloride. 

The  Air.  —  As  already  stated  above,  the  air  consists  of  about 
one  fifth  oxygen  and  four  fifths  nitrogen  by  volume.  That 
these  gases  are  not  chemically  bound  to  each  other  but  simply 
mixed  is  evident  from  the  following  facts :  (1)  When  the  air 
is  cooled,  the  oxygen  condenses  to  a  liquid  first,  leaving  the 
nitrogen  in  form  of  a  gas ;  or  when  liquid  air  is  boiled,  the 
nitrogen  distills  off  first,  leaving  nearly  pure  liquid  oxygen 
behind.  (2)  The  composition  of  the  air,  though  nearly  con- 
stant, varies  somewhat  at  different  times  and  places,  the  oxy- 
gen content  commonly  varying  from  20.9  to  21.0  per  cent. 
(3)  Water  will  dissolve  air  to  some  extent.  When  the  water 
is  then  deprived  of  this  air  by  boiling,  the  air  expelled  from  the 
water  is  richer  in  oxygen  and  poorer  in  nitrogen  than  ordinary 
air.  Thus  in  air  expelled  from  water  the  oxygen  content  is 
35.1  per  cent  and  the  nitrogen  is  64.9  per  cent;  whereas  in 
ordinary  air  the  corresponding  figures  are  20.96  and  79.04  per 
cent,  respectively.  (4)  Air  made  by  mixing  four  volumes  of 
nitrogen  and  one  of  oxygen  behaves  like  ordinary  air.  During 


144  OUTLINES   OF   CHEMISTRY 

the  preparation  of  the  mixture  there  is  neither  a  change  of 
volume  nor  of  temperature. 

The  amount  of  oxygen  and  nitrogen  in  the  air  may  be  deter- 
mined by  passing  air  freed  from  carbon  dioxide  and  moisture 
over  red-hot  copper  and  collecting  and  weighing  the  nitrogen, 
which  is  not  absorbed  by  the  copper.  The  oxygen  is  determined 
by  the  increase  of  weight  of  the  copper,  which  has  united  with 
the  oxygen  of  the  air  passed  over  it.  This  is  the  method  em- 
ployed by  Dumas  and  Boussingault  in  1841.  Another  method 
consists  of  mixing  a  carefully  measured  volume  of  air  with  a 
known  excess  of  hydrogen  and  exploding  the  mixture  by  means 
of  an  electric  spark.  In  this  way  the  oxygen  completely  unites 
with  hydrogen  to  form  water  whose  volume  is  extremely  small 
relatively.  And  so  from  the  diminution  of  the  gaseous  volume 
after  the  explosion  and  the  known  relation  of  the  volumes  of 
hydrogen  and  oxygen  in  water,  the  amount  of  oxygen  in  the  air 
may  readily  be  computed. 

As  a  result  of  the  average  of  numerous  analyses  of  air,  it  has 
been  found  that  the  atmosphere  consists  essentially  of  21  volumes  of 
oxygen  to  79  volumes  of  nitrogen,  or  of 23. 2  per  cent  oxygen  and  76.8 
per  cent  nitrogen  by  weight.  Usually  the  composition  of  differ- 
ent samples  of  air  varies  from  these  figures  by  only  one-tenth  of 
a  per  cent.  A  liter  of  air  at  0°  and  760  mm.  pressure  weighs 
1.2933  grams.  That  the  ratio  of  oxygen  to  nitrogen  in  air  is  so 
nearly  constant  is  due  to  the  fact  that  while  animals  are  con- 
tinually using  up  oxygen  in  respiration,  plants  are  on  the  other 
hand  giving  off  oxygen  to  the  air.  Furthermore,  the  atmos- 
phere is  so  vast  that  the  ordinary  processes  of  combustion  make 
scarcely  a  preceptible  impression  upon  its  oxygen  content. 

Besides  oxygen  and  nitrogen,  the  air  always  contains  water 
vapor,  ammonia,  hydrogen,  nitric  acid,  carbon  dioxide,  dust 
particles  of  organic  as  well  as  inorganic  nature,  and  various 
bacteria  and  other  microbes.  All  of  these  constituents  are,  how- 
ever, quite  variable  in  amount.  In  the  neighborhood  of  cities, 
sulphur  dioxide  and  hydrocarbon  gases  have  also  been  found  in 
the  air.  The  amount  of  water  vapor  in  the  air  varies  greatly 
with  the  locality  and  the  temperature.  Air  saturated  with 
moisture  at  0°  contains  4.87  grams  of  water  vapor  per  cubic 
meter,  while  at  20°  it  contains  17.157  grams.  As  stated  in 
connection  with  the  consideration  of  water,  the  air  is  usually 


NITROGEN,   AIR,   AND   THE   HELIUM   GROUP  145 

saturated  to  only  about  two  thirds  of  its  capacity*  The  amount 
of  moisture  in  the  air  is  best  found  by  passing  a  given  volume 
of  it  through  sulphuric  acid  and  phosphorus  pentoxide  and 
determining  the  increase  in  weight  of  these  drying  agents. 

In  normal  country  air  or  air  over  the  sea,  there  are  about 
3  volumes  of  carbon  dioxide  in  every  10,000  volumes.  In  city 
air,  the  carbon  dioxide  content  is  often  from  6  to  7  volumes 
per  10,000.  In  closed  rooms  where  the  air  is  contaminated  by 
respiration  and  combustion  of  illuminating  gas  or  oil,  the  carbon 
dioxide  content  may  run  as  high  as  6  to  8  times  the  latter  amount. 
Air  containing  more  than  7  volumes  of  carbon  dioxide  is  con- 
sidered harmful  for  continuous  breathing.  The  carbon  dioxide 
in  the  air  is  determined  by  passing  a  known  volume  of  the  latter 
through  baryta  water  and  weighing  the  barium  carbonate  formed. 
The  reaction  that  takes  place  is : 

Ba(OH)2  +  C02  =BaC03+  H20. 

City  air  contains  more  carbon  dioxide  than  country  air,  mainly 
because  of  large  amounts  of  fuel  consumed  in  cities,  and  because 
in  the  country  the  carbon  dioxide  is  taken  from  the  air  to  a  con- 
siderable extent  by  plants. 

Ammonia  occurs  in  the  air  in  very  minute  and  variable 
amounts  hardly  exceeding  from  0.5  to  1  gram  per  10,000  grams 
of  air.  It  arises  as  a  decomposition  product  of  organic  matter 
and  is  not  present  in  the  air  in  the  free  state,  but  is  commonly 
combined  with  nitrous  and  nitric  acids  as  nitrites  and  nitrates. 
The  latter  acids  are  formed,  as  already  mentioned,  when  light- 
ning discharges  in  the  air.  The  ammonium  salts  are  washed 
from  the  atmosphere  during  rains.  Thus  they  get  into  the  soil 
and  serve  as  an  important  nitrogen  supply  for  plants.  The 
latter  get  their  nitrogen  from  this  source  or  from  manures. 
Leguminous  plants,  like  peas,  beans,  and  the  various  varieties 
of  clover,  are  able  to  get  their  nitrogen  supply  from  little 
nodules  which  are  produced  on  their  roots  by  certain  species  of 
bacteria,  which  get  nitrogen  directly  from  the  atmosphere  that 
circulates  in  the  porous  soil.  These  nodules  may  contain  up  to 
five  per  cent  of  nitrogen.  Many  plants  are  incapable  of  assimi- 
lating nitrogen  in  form  of  ammonia.  The  latter  must  first  be 
oxidized.  This  is  brought  about  by  bacterial  action  in  the  soil. 
The  amount  of  nitric  acid  in  the  air  is  small  and  very  variable. 

L 


146  OUTLINES   OF   CHEMISTRY 

Rain  water  has  been  observed  to  contain  0.14  part  of  nitrogen 
as  nitrates  per  million  parts  of  water  on  the  average  in  some 
localities. 

As  has  already  been  stated,  it  is  doubtful  whether  ozone  is 
normally  present  in  the  air.  The  effects  observed  on  starch 
potassium  iodide  paper  may  well  be  due  to  hydrogen  peroxide 
or  higher  oxides  of  nitrogen. 

The  hydrogen  content  of  the  air  varies  considerably.  Ray- 
leigh  found  it  to  be  0.003  per  cent  by  volume.  Dewar  isolated 
0.001  per  cent  hydrogen  from  liquid  air ;  while  Gautier  claims 
to  have  found  as  much  as  0.02  per  cent.  The  hydrogen  gets 
into  the  atmosphere  from  volcanic  gases,  and  as  a  product  of 
bacterial  action.  During  the  process  of  the  decay  of  animal 
and  vegetable  matter  there  are  also  still  other  gases  produced, 
which  enter  the  atmosphere.  These  are,  however,  soon  oxi- 
dized, especially  in  the  presence  of  sunlight. 

The  particles  of  solid  matter  in  the  air  frequently  carry 
bacteria.  As  a  rule  the  bacteria  found  in  the  air  are  harmless, 
though  pathogenic  organisms  do  get  into  the  air,  especially  in 
the  sick  room  and  in  crowded  cities.  Dry  weather  and  winds 
increase  the  amount  of  dust  in  the  air,  and  also  the  number  of 
organisms  that  cling  to  dust  particles.  The  spores  of  molds 
and  microbes  producing  fermentation  and  putrefaction  are 
practically  always  present  in  the  air.  By  filtering  the  latter 
through  plugs  of  cotton,  dust  and  microbes  may  be  removed 
from  the  air.  Normal  air  contains  but  4  or  5  microorganisms  per 
liter.  The  waters  of  rivers  and  inland  lakes  contain  from  5,000 
to  20,000  organisms  per  cubic  centimeter,  whereas  the  soil  con- 
tains about  5  times  the  latter  number  per  cubic  centimeter. 
Thus,  it  is  clear  that  the  air  is  relatively  free  from  organisms. 
The  latter  get  into  the  air  chiefly  from  the  dry  soil,  or  dry  ob- 
jects, as  dust  is  carried  from  them  by  currents  of  air.  Dust 
particles  act  as  nuclei  for  the  condensation  of  moisture  in  the 
formation  of  fogs. 

The  air  that  is  exhaled  by  animals  and  human  beings  con- 
tains, besides  carbon  dioxide,  organic  material ;  arid  it  is  chiefly 
the  latter  which  gives  rise  to  headache  and  general  depression 
that  one  experiences  in  crowded  rooms.  The  decomposition 
products  of  this  organic  matter  give  rise  to  unpleasant  odors 
which  are  frequently  met  in  crowded,  poorly  ventilated  rooms. 


NITROGEN,   AIR,   AND   THE   HELIUM   GROUP  147 

Liquid  air  is  .now  produced  on  a  commercial  scale.  The 
methods  employed  are  founded  upon  the  principle  that  by  sub- 
jecting a  gas  to  very  high  pressure  and  then  allowing  it  to 
escape  through  a  small  orifice,  the  remaining  gas  is  cooled,  due 
to  the  heat  absorbed  in  expansion.  Thus,  air  compressed  to 
about  200  atmospheres  (i.e.  3000  pounds  to  the  square  inch)  is 
cooled  to  room  temperature  by  means  of  cold  water,  and  this  air 
is  then  allowed  to  escape  from  the  long  tube  in  which  it  is  con- 
tained, through  an  orifice  the  size  of  which  is  controlled  by 
means  of  a  needle  valve.  The  air  thus  enters  another  chamber 
which  surrounds  the  first  tube.  The  outflow  is  regulated  so 
that  in  this  second  chamber  the  pressure  of  the  air  is  about  20 
atmospheres.  In  thus  coursing  from  the  first  chamber  into  the 
second  against  a  pressure  of  20  atmospheres  work  is  done,  and 
the  heat  required  to  do  this  work  is  taken  from  the  tube  contain- 
ing the  highly  compressed  air.  The  apparatus  is  carefully  in- 
sulated from  the  surroundings  by  means  of  wool.  After  thus 
continuing  to  feed  the  apparatus  compressed  air  for  a  few 
hours,  the  temperature  in  the  inner  tube  becomes  so  low  that  the 
air  liquefies  and  can  then  be  drawn  off.  It  is  turbid  in  appear- 
ance, due  to  the  solid  particles  of  carbon  dioxide  and  water  that 
it  contains.  These  may  be  filtered  off.  The  filtrate  is  a  clear 
liquid  of  bluish  hue.  Liquid  air  rapidly  changes  its  compo- 
sition, since  nitrogen  evaporates  faster  than  oxygen.  Liquid 
air  boils  at  about  — 190°.  The  boiling  point  of  nitrogen  is 
— 190.5°  and  that  of  oxygen  is  —  182.5°.  After  a  time,  nearly 
all  the  nitrogen  has  evaporated,  leaving  practically  only  oxygen 
behind.  The  latter  is  put  on  the  market  in  steel  cylinders 
as  compressed  oxygen. 

The  Elements  of  the  Helium  Group.  —  It  was  found  by  Lord  Ray- 
leigh  that  a  liter  of  nitrogen  prepared  from  air  weighs  1.2572 
grams  and  that  the  same  volume  of  nitrogen  prepared  from  chem- 
ical compounds  weighs  1.2521  grams.  This  led  Rayleigh  and 
Ramsay  to  investigate  the  composition  of  air  more  carefully,  with 
the  result  that  they  discovered  in  it  the  new  element  argon  in 
1894.  Argon  may  be  prepared  by  passing  air  over  heated  copper 
to  take  out  the  oxygen,  and  then  over  hot  magnesium  or  lithium 
to  absorb  the  nitrogen.  Or  air  may  be  mixed  with  an  excess  of 
oxygen  and  subjected  to  the  action  of  the  electric  spark,  the  gas 
being  kept  over  caustic  potash  solution  to  absorb  the  oxides  of 


148  OUTLINES   OF   CHEMISTRY 

nitrogen  formed.  In  the  latter  method,  the  excess  of  oxygen 
may  finally  be  absorbed  by  passing  the  gas  over  heated  copper 
or  by  treatment  with  an  alkaline  solution  of  pyrogallic  acid. 
About  0.9  per  cent  of  the  air,  by  volume,  consists  of  argon. 
The  gas  has  properties  similar  to  those  of  nitrogen ;  but  argon 
has  thus  far  not  been  obtained  in  combination  with  other  ele- 
ments. It  is  very  inert,  chemically,  whence  its  name  argon, 
meaning  inactive.  The  boiling  point  of  argon  is  — 186°  and 
the  melting  point  — 189°.  It  is  more  soluble  in  water  than  nitro- 
gen. At  room  temperatures  about  40  cc.  of  argon  are  dissolved 
by  1  liter  of  water.  The  gas  has  consequently  been  found  in  all 
natural  waters.  Argon  is  19.95  times  as  heavy  as  hydrogen  ;  its 
molecular  weight  is  consequently  39.9.  As  it  combines  with  no 
other  elements  whatever,  its  atomic  weight  cannot  be  ascertained 
by  the  usual  means.  The  molecular  heat  of  gases  containing 
two  atoms  to  the  molecule  is  approximately  5  Cal.  ;  in  the  case 
of  mercury  vapor,  which  contains  but  one  atom  in  the  molecule, 
the  molecular  heat  is  but  2.5  Cal.  Now  it  has  been  found  that 
to  heat  39.9  grams  of  argon  one  degree  requires  2.5  Cal.,  con- 
sequently argon,  like  mercury,  contains  but  one  atom  in  its 
molecule.  The  molecular  weight  and  atomic  weight  of  argon 
are  consequently  the  same,  namely  39.9.  That  argon  is  a  sim- 
ple substance  is  supported  by  the  fact  that  it  has  a  constant 
boiling  point,  and  that  by  shaking  the  gas  with  water  the  dis- 
solved portion  is  identical  with  the  undissolved  portion.  Argon 
is  extraordinarily  stable,  and  since  it  has  not  been  decomposed 
into  anything  simpler,  it  must  be  regarded  as  an  element. 

Helium,  neon,  krypton,  and  xenon,  four  additional  new  ele- 
ments, were  later  also  discovered  in  the  air  by  Ramsay  and 
Travers. 

Helium  was  known  to  exist  in  the  sun,  whence  its  name.  In 
1895  Ramsay  prepared  helium  by  heating  the  mineral  cleveite 
with  sulphuric  acid.  The  element  exhibits  a  characteristic 
yellow  line  in  its  spectrum.  This  line  had  previously  been 
observed  in  the  spectrum  of  the  sun  by  Lockyer,  who  ascribed 
it  to  an  element  which  he  called  helium,  then  unknown  on  the 
earth.  Helium  does  not  unite  with  any  other  element.  Its 
molecular  weight  is  3.99  and  its  atomic  weight  is  the  same. 
About  1.4  cc.  of  helium  are  dissolved  by  100  cc.  of  water  at 
room  temperature.  Helium  was  liquefied  in  1908  by  Professor 


NITROGEN,   AIR,   AND   THE   HELIUM   GROUP  149 

Kammerlingh  Onnes  of  the  University  of  Leiden.  Its  boiling 
point  is  —  268.7°,  and  its  specific  gravity  in  the  liquid  state  is 
0.15.  This  gas  is  the  most  difficult  one  to  liquefy,  and  for  sev- 
eral years  it  had  resisted  all  attempts  to  condense  it  to  a  liquid. 
We  may  now  say  that  all  known  gases  have  been  liquefied.  In 
the  air  helium  occurs  to  the  extent  of  1  to  2  volumes  per  million 
volumes. 

Ramsay  found  neon,  krypton,  and  xenon  in  the  argon  pre- 
pared from  air.  Helium  and  neon  are  dissolved  in  the  liquid 
argon,  from  which  they  are  expelled,  together  with  argon,  as 
the  temperature  rises.  The  residue  Ramsay  subjected  to  further 
fractional  distillation,  and  so  separated  krypton  and  xenon  from 
each  other.  By  cooling  a  mixture  of  neon  and  helium  with 
liquid  hydrogen,  neon  solidifies,  while  the  helium  remains  in 
the  gaseous  state.  Neon,  krypton,  and  xenon  are  inert  gases 
that  combine  with  no  other  elements;  they  are  mono-atomic. 
Their  atomic  weights  as  determined  from  their  densities 
are  :  — 

Neon         .         .         .        20.2 

Krypton    .         .         .         82.92 

Xenon       .         .         .       130.2 

Their  boiling  points  are  as  follows :  — 

Neon          ...        —  243°  (approximately) 
Krypton    .  - 152° 

Xenon       .         .         .       -109° 

Krypton  melts  at  -  169°  and  xenon  at  -  140°. 

The  following  table  gives  the  amounts  of  the  gases  of  the 
helium  group  contained  in  one  cubic  meter,  i.e.  1000  liters,  of 
air :  — 

Helium  .  .  0.0015      liter  .  .  =    0.0002T  gram 

Neon  .  .  0.015        liter  .  .  =    0.01339  gram 

Argon  .  .  9.4            liters  .  .  =16.76         grams 

Krypton  .  .  0.00005    liter  .  .  =    0.00018  gram 

Xenon  .  .  0.000006  liter  .  .  =    0.00003  gram 


CHAPTER  XI 

COMPOUNDS   OF  NITROGEN   WITH   HYDROGEN   AND   WITH 
THE   HALOGENS 

History  and  Occurrence  of  Ammonia.  — Ammonia  is  by  far 
the  most  important  compound  of  nitrogen  with  hydrogen.  Up 
to  1774  it  was  known  only  in  its  aqueous  solution,  which  Glauber 
called  "  spiritus  volatilis  salis  armoniaci,"  and  which  was  later 
named  spirits  of  hartshorn  and  spirits  of  sal  ammoniac.  Am- 
monia was  prepared  by  treating  sal  ammoniac,  that  is,  ammo- 
nium chloride,  with  lime  or  some  other  alkali,  whence  the  name 
spirits  of  sal  ammoniac.  It  may  also  be  obtained  by  heating 
hoofs  and  horns  of  animals  out  of  contact  with  the  air,  whence 
the  term  spirits  of  hartshorn.  The  process  of  thus  heating 
substances  out  of  contact  of  the  air  in  a  retort  and  decompos- 
ing them  into  other  products  is  called  destructive  distillation. 
Priestley  discovered  ammonia  gas  in  1774  by  evolving  it  from 
lime  and  ammonium  chloride  and  collecting  it  over  mercury. 
He  called  it  "  alkaline  air  " ;  for  the  gas  turns  red  litmus  blue 
and  acts  in  other  ways  like  a  strong  alkali.  As  stated  in  con- 
nection with  nitrogen,  ammonia  occurs  in  small  amounts  in  the 
atmosphere  in  form  of  ammonium  salts,  particularly  as  ammo- 
nium carbonate.  It  is  a  product  of  the  decomposition  of  all 
vegetable  and  animal  matter,  and  hence  is  found  in  all  natural 
waters  and  soils.  In  the  form  of  salts,  mainly  nitrate  and 
nitrite,  it  occurs  in  rain  water.  Its  occurrence  in  soils  is  im- 
portant, for  it  is  a  fertilizer.  In  the  neighborhood  of  volcanoes, 
particularly  those  of  Tuscany,  ammonia  occurs  in  the  form  of 
sulphate  and  chloride  of  ammonium.  Ammonium  chloride  used 
to  be  prepared  in  Egypt  in  the  oasis  near  the  temple  of  Jupiter 
Ammon,  from  the  soot  obtained  by  heating  camel's  dung  which 
was  used  as  fuel;  thus  the  salt  received  its  name  sal  ammoniac, 
from  which  comes  the  term  ammonia. 

Preparation  and  Properties  of  Ammonia.  —  When  a  mixture 
of  nitrogen  and  hydrogen  is  subjected  to  the  silent  electrical 

150 


AMMONIA   AND   OTHER   NITROGEN   COMPOUNDS  151 

discharge  as  in  making  ozone,  small  amounts  of  ammonia  are 
formed  by  direct  union  of  the  elements,  thus  :  — 


The  common  method  of  preparing  ammonia  is  by  heating  am- 
monium chloride  with  slaked  lime  :  — 

2  NH4C1  +  Ca(OH)2  =  CaCl2  +  2  H2O  +  2  NH3. 

Any  other  ammonium  salt  may  be  used  instead  of  the  chloride, 
and  other  alkalies,  like  sodium  or  potassium  hydroxide,  may 
serve  in  place  of  lime,  which,  however,  is  the  cheapest. 

By  the  reduction  of  nitrates  and  nitrites  with  nascent  hydro- 
gen, ammonia  may  be  produced,  thus  :  — 

KN03  +  8  H  =  KOH  +  2  H2O  +  NH3. 
KN02  +  6  H  =  KOH  +  H2O  +  NH8. 

By  the  dry  distillation  of  nitrogenous  animal  and  vegetable 
material  ammonia  is  formed.  So  by  heating  coal  (which  rep- 
resents the  remains  of  vegetation  of  the  carboniferous  age)  out 
of  contact  with  air,  as  is  done  in  the  manufacture  of  illuminat- 
ing gas  from  coal,  ammonia  is  produced.  The  coal  gas  formed 
is  passed  through  water,  which  readily  dissolves  the  ammonia, 
and  it  is  from  these  ammoniacal  liquors  of  the  gas  works  that 
the  ammonia  of  commerce  is  almost  entirely  obtained  at  present. 
From  these  liquors  the  gas  is  expelled  by  heating  with  slaked 
lime.  The  ammonia  so  liberated  is  passed  into  sulphuric  acid, 
and  the  sulphate  of  ammonium  thus  formed  is  purified  by  re- 
erystallization.  From  this  pure  salt,  pure  ammonia  and  other 
ammonium  products  are  in  turn  prepared. 

By  heating  organic  nitrogenous  products  with  strong  alkalies, 
ammonia  is  produced.  This  is  frequently  used  in  ascertaining 
the  amount  of  nitrogen  in  organic  substances,  particularly  as 
they  occur  in  fertilizers,  sewage,  drinking  water,  etc.  In  this 
process  a  strong  solution  of  caustic  potash  and  potassium  per- 
manganate is  frequently  employed.  The  latter  substance  is  a 
powerful  oxidizing  agent  and  so  aids  in  the  destruction  of  the 
organic  material.  Animal  matter  when  heated  with  fuming 
sulphuric  acid  is  decomposed,  the  nitrogen  being  converted 
into  ammonia,  which  unites  with  the  sulphuric  acid,  forming 
ammonium  sulphate.  This  process  (known  as  Kjeldahl's 


152  OUTLINES   OF   CHEMISTRY 

method)  is  of  importance  in  the  chemical  analysis  of  nitroge- 
nous organic  substances. 

Ammonia  is  a  colorless  gas  of  a  strong,  peculiar,  penetrating 
odor.  It  is  0.59  time  as  heavy  as  air.  It  may  be  condensed 
to  a  liquid  which  boils  at  —  32.5°.  It  has  also  been  obtained  in 
form  of  white  crystals  that  melt  at  —78°.  The  specific  gravity 
of  liquid  ammonia,  taken  under  pressure  at  0°,  Is  0.6233.  In 
water  the  gas  is  extremely  soluble  At  0°,  1  volume 
of  water  absorbs  1148  volumes  of  ammonia,  while  at 
16°  and  50°,  only  764  volumes  and  306  volumes,  respec- 
tively, are  absorbed.  On  boiling  an  aqueous  solution 
of  ammonia,  the  gas  is  completely  expelled,  which  fact 
is  frequently  used  in  laboratories  for  preparing  ammonia 
gas.  On  account  of  its  solubility  in  water,  ammonia 
gas  is  collected  over  mercury,  or  simply  by  displace- 
ment of  air,  the  vessel  in  which  gas  is  to  be  collected 
being  supported  with  the  bottom  upward,  for  the  gas  is 
but  little  more  than  half  as  heavy  as  air. 

The  weight  of  a  liter  of  ammonia  gas  at  0°  and 
760  mm.  is  0.7635  gram,  and  since  the  gas  consists  of 
82.27  per  cent  nitrogen  and  17.73  per  cent  hydrogen 
by  weight,  its  formula  is  NH3.  By  electrolyzing  an 
aqueous  ammonia  solution,  to  which  some  common  salt 
has  been  added  to  make  the  solution  conduct  better, 
three  volumes  of  hydrogen  are  obtained  to  one  volume  of 
nitrogen.  The  apparatus  used  for  this  purpose  is  the 
same  as  that  shown  in  Fig.  2. 

Again,  when  a  given  volume  of  ammonia  gas  (Fig.  36) 
is  treated  with  a  concentrated  solution  of  potassium 
hypobromite,  nitrogen  is  formed  which  occupies  half 
the  volume  of  the  original  ammonia.   Care  must  be  taken 
not  to  admit  air  irito  the  tube  during  the  experiment. 
We  thus  see  that  2  volumes  of  ammonia  yield  1  volume 
FIG"  36     °f  nitrogen,  while  by  electrolysis  3  volumes  of  hydrogen  and 
1  volume  of  nitrogen  were  obtained  from  ammonia.     Con- 
sequently, these  volume  relations  may  be  expressed  thus :  — 

3  volumes  hydrogen  + 1  volume  nitrogen  =  2  volumes  ammonia  gas. 

We   have   here   another   excellent  confirmation   of  the  law 
of  Gay-Lussac  of  the  combination  of  gases  by  volume.     By 


AMMONIA  AND   OTHER   NITROGEN   COMPOUNDS 


153 


Avogadro's  hypothesis,  equal  volumes  of  gases  contain  an  equal 
ftumber  of  molecules,  hence  :  — 

3  molecules  hydrogen  +  1  molecule  nitrogen  =  2  molecules  ammonia 
or  3  H2  +  N2  =  2  NH3. 

While  it  is  true  that  a  mixture  of  3  volumes  of  hydrogen 
and  1  volume  of  nitrogen  when  subjected  to  the  electric  spark 
yields  small  amounts  of  ammonia,  it  is  also  the  case  that  when 
the  latter  gas  is  thus  treated  it  is  partly  decomposed  into  nitro- 
gen and  hydrogen.  The  reaction  is  thus  a  reversible  one  :  — 


If  none  of  the  gases  are  removed,  an  equilibrium  is  finally 

slowly  reached,  which  is  the  same  in  each  case,  the  gases  con- 

sisting of  2  per  cent  ammonia  and  98 

per  cent  of  uncombined  nitrogen  and 

hydrogen.      So  if  ammonia  gas  con- 

tained in  the  closed  limb  of  the  appa- 

ratus shown  in  Fig.  37  is  treated  with 

the  electric  spark,  the  volume  of  hy- 

drogen plus  nitrogen  formed  will  be 

nearly  twice  that  of  the  original  volume 

of  ammonia.    If,  however,  the  ammonia 

is    removed    (absorbed    by    sulphuric 

acid,    for   example)    as    fast   as   it   is 

formed,  the  reaction  completes   itself 

from  left  to  right  as  would  be  expected, 

according  to  the  law  of  mass  action. 

When  ammonia  is  oxidized,  as,  for 
instance,  by  passing  it  over  hot  copper 
oxide,  the  latter  is  reduced  and  the 
only  products  formed  are  water  and 
nitrogen,  thus  showing  that  the  gas  is 
composed  of  hydrogen  and  nitrogen. 
By  thus  oxidizing  a  definite  volume  of 
ammonia  and  weighing  the  water  and  nitrogen  formed,  the 
percentage  composition  of  ammonia  has  been  determined. 
The  results  of  such  analyses  have  already  been  given  above. 
On  account  of  its  hydrogen  content  ammonia  will  burn. 
The  action  proceeds  in  oxygen,  but  not  in  air.  Thus,  when 


FIG.  37. 


154 


OUTLINES   OF   CHEMISTRY 


in  a   flask   (Fig.    38)   strong   ammonia  water    is   heated    till 
ammonia  is  copiously  evolved,  and  oxygen  is  then  conducted 

into  the  gas,  the  mixture  when 
lighted  will  burn  at  the  mouth  of 
the  flask.  The  products  are  mainly 
water  and  nitrogen,  though  ni- 
trous and  nitric  acids  are  also 
formed  to  a  slight  extent.  These 
acids  unite  with  the  excess  of 
ammonia  to  form  ammonium 
nitrite  and  nitrate.  The  latter 
salts  are  more  copiously  formed 
when  a  heated  spiral  of  platinum 
wire  (Fig.  39)  is  hung  into  a 
mixture  of  oxygen  and  ammonia 
*  FIG.  38.  gases.  The  platinum  continues 

to  glow,  and  white  fumes  form  which  consist  of  the  salts  men- 
tioned.    The  platinum  here  acts  as  a  catalytic  agent. 

Ammonia  water  is  lighter  than  water.  The  saturated  solu- 
tion at  14°  C.  contains  36  per  cent  NH3  and  has  a  specific  gravity 
of  0.8844.  It  is  sold  as  a  con- 
centrated ammonia,  and  may  be 
diluted  to  any  other  strength 
desired. 

Ammonia   unites  directly  with 
acids,  forming  salts,  thus :  — 

NH3  +  HC1  =  NPI4C1. 
2  NH3  +  H2SO4  =  (NH4)2S04. 


NH3  +  HNO 


NH4NO3. 


We  may  regard  these   salts   as 

derived  from  the   acids   by  the 

replacement   of    each   hydrogen 

atom  by  the   group   NH4.       So 

we  may  also  consider   that   the 

group  NH4  plays  the  role  of  an 

atom  of  a  univalent  metal,  like 

Na  or  K.      For  this  reason,  NH4 

is  called  ammonium,  the  ending  um  being  used  to  indicate  that 

chemically  it  is  analogous  to  a  metal.     When  ammonia   dis- 


AMMONIA   AND   OTHER  NITROGEN  COMPOUNDS  155 

solves  in  water,  much  heat  is  evolved,  and  we  may  consider 
that  the  addition  product  formed  is  NH4OH,  thus  :  — 

NH3+H20  =  NH4OH. 

The  latter  has  not  been  isolated  ;  but  the  aqueous  solutions  act 
as  though  this  substance  were  contained  in  them.  So,  for 
example,  when  ammonia  water  is  neutralized  by  hydrochloric 
or  sulphuric  acid,  the  action  may  be  expressed  thus  :  — 

NH4OH  +  HC1=  NH4C1  +  H2O. 
2  NH4OH  +  H2S04  =  (NH4)2S04  +  2  H2O. 

At  dull  red  heat  all  ammonium  salts  are  volatilized.  This  is 
a  very  important  fact  in  chemical  analysis.  In  many  cases  the 
salts  are  simply  broken  up  into  ammonia  and  the  free  acid, 
thus  :  — 


In  such  instances,  which  are  typical  cases  of  dissociation,  the 
actions  are  reversible,  the  products  again  uniting  as  the  tem- 
perature is  lowered.  The  vapor  of  ammonium  chloride  con- 
tains all  three  products  named  in  the  above  equation  ;  which 
was  demonstrated  by  Henri  Saint  Claire  Deville,  who  separated 
hydrochloric  acid  and  ammonia  from  these  vapors  by  diffusion, 
making  use  of  the  fact  that  ammonia,  the  lighter  gas,  diffuses 
more  rapidly  than  hydrochloric  acid. 

By  means  of  chlorine,  ammonia  is  decomposed  :  — 

2  NH3  +  3  C12  =  6  HC1  +  N2. 

When  excess  of  ammonia  is  present,  the  latter  at  once  unites 
with  the  hydrochloric  acid  formed  and  produces  ammonium 
chloride. 

Ammonia  acts  on  many  metals.     Thus,  sodium  and  potassium 
when  heated  act  on  ammonia  as  follows  :  — 

2  Na  4-  2  NH3  =  2  NaNH2  +  Ha. 
2  K  +  2  NH3  =  2  KNH2  +  H2. 

The  compounds  formed  are  sodium  amide  and  potassium  amide. 
They  are  decomposed  by  water  into  ammonia  and  the  hydrox- 
ide of  the  metal,  so,  for  instance  :  —  . 

KNH,  +  H20  =  KOH  +  NHa. 


156  OUTLINES   OF   CHEMISTRY 

The  nitrides  of  lithium  NLi3  and  of  magnesium  N2Mg3  may 
be  regarded  as  ammonia  in  which  the  hydrogen  atoms  are  re- 
placed by  the  respective  metals.  These  nitrides  may  be  formed 
by  igniting  the  metals  in  ammonia.  On  treating  the  nitrides 
with  water,  ammonia  and  the  metallic  hydroxides  are  formed, 
thus  :  — 

NLi3  +  3  H2O  =  3  LiOH  +  NH3. 

6  H20  =  3  Mg(OH)2  +  2  NH3. 


The  fact  that  ammonia  water  will  dissolve  metals  like  zinc, 
copper,  and  silver  and  also  their  oxides,  is  used  in  cleaning 
tarnished  metallic  articles,  for  the  action  of  ammonia  is  not  as 
drastic  as  that  of  an  acid,  and,  furthermore,  the  ammonia  readily 
evaporates  after  use. 

Liquid  ammonia  is  a  great  solvent.  In  this  respect  it  is 
similar  to  water;  for  it  dissolves  many  salts,  forming  with  some 
of  them  addition  products  that  are  analogous  to  hydrates  that 
water  forms  with  salts.  Thus,  with  water  copper  sulphate 
forms  the  compound  CuSO4  •  5  H2O,  in  which  the  water  is  spoken 
of  as  water  of  crystallization.  Similarly  with  dry  ammonia 
copper  sulphate  forms  the  compound  CuSO4-  5  NH3,  in  which 
the  ammonia  may  be  called  ammonia  of  crystallization.  The 
properties  of  liquid  ammonia  solutions  have  been  investigated 
by  E.  C.  Franklin  in  recent  years. 

Liquid  ammonia  is  much  like  water  in  that  it  has  a  high 
specific  heat,  1.02  between  0°  and  20°,  and  a  high  latent  heat 
of  vaporization.  It  takes  316  cal.  to  vaporize  1  gram  of  liquid 
ammonia  at  0°.  This  fact  is  used  in  the  artificial  ice  machines, 
in  which  liquid  ammonia  is  evaporated  in  tubes  which  are  sur- 
rounded with  concentrated  calcium  chloride  solution.  The 
evaporation  of  the  ammonia  requires  heat,  which  is  abstracted 
from  the  brine;  and  this  cold  brine  is  then  distributed  by  means 
of  a  system  of  tubes  to  the  places  where  refrigeration  is  re- 
quired. The  ammonia  is  again  liquefied  by  compression  with  a 
powerful  pump,  and  so  it  can  be  used  over  and  over  again  in  a 
closed  system  of  tubes.  The  calcium  chloride  brine  also  circu- 
lates in  a  closed  system  of  tubes,  the  brine  after  becoming 
warmed  being  returned  and  again  chilled.  That  a  considerable 
lowering  of  temperature  can  be  produced  by  the  evaporation  of 
ammonia  may  be  shown  by  a  simple  experiment.  When  a  con- 


AMMONIA   AND   OTHER  NITROGEN   COMPOUNDS  157 

centrated  aqueous  solution  is  placed  in  a  flask  standing 
upon  a  board  that  is  wet  with  a  few  cubic  centimeters 
of  water,  and  a  strong  current  of  air  is  passed  into  the 
solution  by  means  of  a  pair  of  bellows,  the  evaporation 
of  the  ammonia  proceeds  so  rapidly  that  in  a  few  minutes 
the  cold  produced  is  sufficient  to  freeze  the  flask  to  the 
board. 

In  ammonia  NH3  nitrogen  is  trivalent,  whereas  in  ammonium 
salts  the  element  is  quinquivalent.  Indeed,  in  most  of  its  com- 
pounds nitrogen  may  be  considered  as  having  a  valence  of 
either  three  or  five. 

Ammonium  salts  are  readily  detected  by  the  fact  that 
ammonia  is  evolved  when  they  are  treated  with  caustic 
alkali.  The  ammonia  gas  is  easily  distinguished  by  its  odor 
and  by  the  fact  that  it  turns  red  litmus  paper  blue. 
When  present  in  very  small  quantities,  as  in  drinking  water, 
ammonium  salts  are  detected  by  means  of  a  solution  of  mer- 
curic iodide  HgI2  in  potassium  iodide.  This  solution  is  made 
strongly  alkaline  by  addition  of  caustic  potash  and  is  then 
known  as  Nesslers  reagent.  When  added  to  a  very  dilute 
solution  of  an  ammonium  salt  a  yellow  color  is  produced. 
In  stronger  solutions  of  ammonium  salts  a  dark  brown  color 
or  a  precipitate  is  formed.  Nessler's  reagent  is  of  great 
importance  in  analyzing  potable  waters,  sewage,  and  the 
like. 

Hydrazine.  —  The  composition  of  hydrazine  or  diamide  is  ex- 
pressed by  the  formula  H2N-NH2.  This  compound  was  dis- 
covered by  Curtius  in  1887.  It  may  be  made  by  the  oxidation 
of  urea, 'thus:  — 

NH2-CO-NH2  +  O  =  H2N-NH2  +  CO2  ; 
or  by  the  reduction  of  hyponitrous  acid,  thus :  — 

H-O-N-N-O-H  +  6  H  =  H2N-NH2  +  2  H2O. 

It  forms  white  crystals  melting  at  —1°.  Liquid  hydrazine 
boils  at  113°.  Its  specific  gravity  at  15°  is  1.013.  It  is 
miscible  with  water  in  all  proportions  and  forms  a  hydrate 
H2N-NH8(OH),  which  melts  at  -  40°  and  boils  at  120°.  Its 


158  OUTLINES  OF  CHEMISTRY 

specific  gravity  is  1.03.  Like  ammonia,  hydrazine  is  a  strong 
base.  Its  solutions  have  a  corrosive  action  on  cork  and  rubber, 
and  even  on  glass,  especially  at  higher  temperatures.  With  acids 
hydrazine  forms  salts.  The  action  is  similar  to  the  formation 
of  ammonium  salts.  A  large  number  of  compounds  derived 
from  hydrazine  by  replacing  one  or  more  of  its  hydrogens  by 
hydrocarbon  radicals  are  known  in  organic  chemistry.  One  of 
these,  namely,  phenyl  hydrazine  (C6H5)HN-NH2,  has  been  of 
special  importance  in  the  synthesis  and  investigation  of  sugars. 
Hydroxylamine.  —  This  compound,  which  was  discovered  by 
Lessen  in  1865,  and  prepared  in  the  pure  state  by  Lobry  de 
Bruyn  in  1891,  may  be  formed  by  the  action  of  nascent  hydro- 
gen either  on  nitric  acid  or  nitric  oxide,  thus  :  — 

HN03  +  6  H  =  2  H20  +  NH2(OH). 
NO  +  3H  =  NH2(OH). 

It  may  be  considered  as  ammonia  NH3  in  which  one  hydrogen 
atom  has  been  replaced  by  the  OH  group.  The  group  NH2  is 
called  the  amido  or  amine  group,  whence  the  name  hydroxyl- 
amine.  It  consists  of  white  hygroscopic  needles  that  melt 
at  33°.  The  boiling  point  is  58°  at  22  mm.  and  70°  at  60  mm. 
pressure.  In  water  hydroxylamine  dissolves  readily.  It  has 
basic  properties,  showing  alkaline  reaction  toward  indicators, 
and  forming  crystalline  salts  with  acids,  thus :  — 

NH2OH  +  HC1  =  NH8(OH)C1. 
2  NH2OH  +  H2S04  =  (NH8OH)2S04. 

These  salts  may  be  regarded  as  ammonium  salts  in  which  one 
hydrogen  has  been  replaced  by  OH.  Hydroxylamine  is,  how- 
ever, a  much  weaker  base  than  ammonia  or  hydrazine.  On 
heating  hydroxylamine  or  its  compounds,  decomposition  sets 
in,  which,  on  account  of  sudden  evolution  of  gas,  may  take 
place  with  explosive  violence.  The  reducing  power  of  hydrox- 
ylamine is  characteristic.  By  treating  a  hot  alkaline  solution 
of  a  cupric  salt  with  hydroxylamine,  red  cuprous  oxide  is  at 
once  formed,  thus  :  — 

4  CuO  +  2  NH2(OH)  =  N2O  +  3  H2O  +  2  Cu2O. 

The  reaction  will  take  place  even  when  hydroxylamine  is 
present  merely  as  1  part  in  100,000. 


AMMONIA  AND   OTHER   NITROGEN    COMPOUNDS  159 

Hydroxylamine  readily  decomposes  into  ammonia,  nitrogen, 
and  water,  thus  :  — 

3  NH2OH  =  NH8  +  N2  +  3  H2O. 

In  organic  chemistry  hydroxylamine  is  of  importance,  because 
with  aldehydes  and  ketones  (which  see)  it  forms  compounds 
known  as  oximes. 

Hydrazoic  Acid.  —  This  compound  has  the  composition  ex- 
pressed by  the  formula  N3H.  It  is  also  called  hydronitric  acid, 
triazoic  acid,  or  azoimide.  It  was  discovered  by  Curtius  in 
1890.  It  may  be  made  by  passing  nitrous  oxide  over  sodium 
amide  at  200°,  and  then  treating  the  resulting  sodium  hydrazo- 
ate  with  dilute  sulphuric  acid.  The  reactions  are  as  follows:  — 

NaNH2  +  N2O  =  H2O  +  NaN3. 
2  NaN8  +  H2S04  =  Na2SO4  +  2  HN3. 

By  carefully  distilling  the  aqueous  solution  produced,  a  solution 
of  the  free  acid  in  water  may  be  obtained.  The  pure  acid  boils 
at  37°.  It  is  a  colorless  liquid  with  a  disagreeable,  penetrating 
odor.  When  inhaled,  it  irritates  the  mucous  membranes.  It 
explodes  with  violence,  forming  nitrogen  and  hydrogen  with 
liberation  of  much  heat,  thus  :  — 

2N3H  =  3N2+H2. 

It  is  a  monobasic  acid,  and  in  this  respect  it  is  similar  to  the 
hydrohalogen  acids.  Its  salts  are  also  unstable  and  liable  to 
explode  with  violence.  It  is  of  interest  to  note  that  the  one 
hydride  of  nitrogen  NH3  is  alkaline  and  the  other  N3H  is  acid. 
The  two  will  combine  to  form  a  salt,  thus :  — 

NH3  +  (N3)H  =  NH4(N8). 

The  empirical  formula  of  NH4(N3),  ammonium  hydrazoate,  is 
N4H4. 

Compounds  of  Nitrogen  with  the  Halogens.  —  With  the  halo- 
gens nitrogen  forms  extremely  unstable  compounds. 

Nitrogen  trichloride  NC13  is  formed  by  the  action  of  chlorine 
upon  ammonium  chloride,  thus  :  — 

NH4C1  +  3  Cla  =  4  HC1  +  NC18. 

The  compound  may  be  prepared  by  the  electrolysis  of  an  aque- 
ous solution  of  ammonium  chloride,  the  chlorine  liberated  acting 


160  OUTLINES   OF   CHEMISTRY 

on  the  solution  according  to  the  above  equation.  Nitrogen 
trichloride  is  a  thin,  yellowish,  oily  liquid  of  specific  gravity 
1.65.  It  is  an  extremely  dangerous  substance  to  deal  with,  for  it 
explodes  with  great  violence  when  heated  or  brought  into  con- 
tact with  substances  like  turpentine  or  phosphorus,  or  when 
exposed  to  sunlight.  Often  the  explosion  occurs  spontaneously, 
which  makes  the  danger  of  working  with  it  very  great  indeed. 
It  has  a  pungent  odor,  and  its  fumes  irritate  the  mucous  mem- 
branes. It  is  soluble  in  hydrocarbons  and  carbon  disulphide, 
the  solutions  thus  formed  being  yellow  in  color  and  compara- 
tively harmless.  At  71°  nitrogen  trichloride  boils  and  may 
be  distilled,  though  the  danger  incurred  in  the  operation  is 
extremely  great.  By  concentrated  hydrochloric  acid  or  aqueous 
ammonia  solution,  nitrogen  trichloride  may  be  decomposed, 

thus  :  — 

NC13  +  4  HC1  =  NH4C1  +  3  C12. 

NC13  +  4  NH4OH  =  3  NH4C1  +  4  H2O  +  Na. 

Nitrogen  trichloride  was  discovered  by  Dulong  in  1811.  In 
working  with  the  substance  he  was  so  unfortunate  as  to  lose 
an  eye  and  three  fingers  in  consequence  of  an  explosion. 

Nitrogen  tribromide  is  a  red,  oily,  highly  explosive  substance 
formed  by  the  action  of  potassium  bromide  on  nitrogen 
chloride.  The  substance  is  believed  to  have  the  composition 
represented  by  the  formula  NBr3. 

Nitrogen  iodide  is  formed  when  iodine  is  treated  with  a  con- 
centrated aqueous  ammonia  solution,  or  when  an  alcoholic  solu- 
tion of  iodine  is  mixed  with  strong  aqueous  ammonia.  The 
compound  is  a  brown  powder  having  the  composition  N2H3T3, 
probably  I3N  =  NH3.  It  is  not  explosive  when  wet  ;  but  when 
dry  it  is  very  explosive,  a  touch  with  a  feather  sufficing  to  cause 
it  to  explode  with  detonation. 

By  treatment  of  silver  hydrazoate  AgN8  with  a  solution  of 
iodine  in  ether,  triazoiodide  IN3  may  be  formed  :  — 


It  is  a  yellow  powder  of  a  very  penetrating  odor,  and  is  ex- 
tremely explosive. 


CHAPTER   XII 

OXY-ACIDS  AND   OXIDES   OP   NITROGEN 

THREE  oxy-acids  and  five  oxides  of  nitrogen  are  known. 
These  are  nitric  acid  HNO3,  nitrous  acid  HNO2,  hyponitrous 
acid  H2N2O2,  nitrogen  pentoxide  or  nitric  anhydride  N2O5, 
nitrogen  peroxide  NO2  or  N2O4,  nitrogen  trioxide  or  nitrous 
anhydride  N2O3,  nitric  oxide  NO,  and  nitrous  oxide  N2O.  In 
the  consideration  of  these  compounds  nitric  acid  will  be  taken 
up  first,  for  from  it  the  other  oxy-acids  and  oxides  named  are 
generally  prepared. 

History,  Occurrence,  and  Preparation  of  Nitric  Acid.  —  Nitric 
acid  was  known  to  the  alchemists  under  the  name  of  aquafortis. 
Up  to  the  seventeenth  century  the  acid  was  prepared  by  heat- 
ing a  mixture  of  saltpeter,  copper  sulphate,  and  alum  according 
to  directions  given  by  the  alchemist  Geber,  who  probably 
lived  in  the  ninth  and  tenth  centuries.  In  this  process  copper 
sulphate  and  alum  yield  sulphuric  acid,  which  unites  with  the 
potassium  of  the  saltpeter,  thus  setting  nitric  acid  free.  Pre- 
pared by  this  method,  the  acid  was  impure.  In  1650  Glauber 
prepared  nitric  acid  by  treating  saltpeter  with  sulphuric  acid, 
and  this  method  is  in  vogue  to  the  present  day.  Though 
Lavoisier  studied  nitric  acid  and  showed  that  it  contained  oxy- 
gen, he  did  not  ascertain  the  real  nature  of  the  acid.  In  1784 
Cavendish  demonstrated  the  nature  of  the  acid  by  preparing  it 
by  passing  an  electric  spark  through  air.  In  this  way  nitrogen 
peroxide  is  formed,  which  in  contact  with  water  yields  nitric 
acid  (see  below). 

It  has  already  been  stated  that  nitric  acid  and  nitrates  occur 
in  small  amounts  in  the  atmosphere.  In  the  soil  and  in  natural 
waters  nitrates  occur  as  the  final  product  of  the  decomposition 
and  oxidation  of  animal  and  vegetable  matter.  The  chief 
source  of  nitric  acid  is  Chili  saltpeter  or  sodium  nitrate. 
Nitric  acid  gets  its  name  from  the  fact  that  it  is  commonly  pre- 
pared from  niter,  saltpeter. 

M  161 


162 


OUTLINES   OF   CHEMISTRY 


By  treating  a  nitrate  like  potassium  or  sodium  nitrate  with 
concentrated  sulphuric  acid,  nitric  acid  is  liberated,  thus  :  — 

NaN03  +  H2S04  =  NaHSO4  +  HNO3. 

In  the  laboratory  the  sodium  nitrate  is  generally  placed  in  a 
glass  retort  (Fig.  40),  sulphuric  acid  is  added,  and  the  mixture 


FIG.  40. 

gently  heated,  when  nitric  acid  distills  over.  On  a  commercial 
scale  Chili  saltpeter  is  treated  with  sulphuric  acid  in  a  cast-iron 
retort,  and  the  nitric  acid  formed  is  condensed  in  bottles  of  stone- 
ware that  contain  a  little  water,  the  last  bottle  being  connected 
with  a  tower  filled  with  coke  over  which  water  trickles  so  as  to 
dissolve  the  acid  vapors  that  still  remain  uncondensed.  Of  late 
stoneware  pipes  are  frequently  employed  instead  of  the  bottles. 
In  this  way  an  aqueous  solution  is  obtained  which  contains 
about  sixty  per  cent  nitric  acid  and  has  a  specific  gravity  of 
1.37.  By  using  dry  sodium  nitrate  and  concentrated  sulphuric 
acid,  the  nitric  acid  obtained  has  a  specific  gravity  of  1.53  and 
is  practically  free  from  water.  On  heating  the  pure  acid,  as  in 
the  process  of  distillation,  it  decomposes  in  part,  thus  :  — 


4  HNO3  =  2  H2O 


4  NO 


O2. 


The  nitrogen  dioxide  forms  reddish  brown  fumes  that  dissolve 
in  the  nitric  acid.     This  solution,  which  fumes  strongly  in  the 


OXY-ACIDS  AND   OXIDES   OF  NITROGEN  163 

air,  is  termed  red  fuming  nitric  acid.  Its  specific  gravity  is 
about  1.54. 

When  the  electric  spark  passes  through  air,  brown  fumes  are 
formed  which  are  nitrogen  dioxide.  These  in  contact  with 
water  form  nitric  acid,  thus :  — 

8  NO2  +  H2O  =  2  HNO3  +  NO. 

This  may  readily  be  shown  by  means  of  the  apparatus  in  Fig. 
35.  Sparks  from  an  induction  coil  are  passed  through  the  air 
between  the  platinum  points  in  the  glass  globe.  After  a  time 
the  gas  in  the  globe  appears  brownish  in  color.  On  shaking 
the  gas  with  water,  and  testing  with  blue  litmus  paper,  the 
presence  of  acid  is  demonstrated.  Many  attempts  have  been 
made  to  use  this  process  for  the  profitable  production  of  nitric 
acid  on  a  commercial  scale.  These  have  been  unsuccessful  till 
recently,  for  the  amount  of  nitric  acid  produced  was  too  small 
as  compared  with  the  electric  power  that  had  to  be  expended, 
even  when  water  power  was  available  for  running  dynamos. 
Of  late,  however,  the  process  has  been  perfected  by  subjecting 
the  electric  arc  formed  to  the  action  of  a  powerful  electro-mag- 
netic field.  In  this  way  arcs  produced  by  means  of  large  alter- 
nating current  dynamos  are  obtained  in  form  of  disks  over  six 
feet  in  diameter,  through  which  air  is  passed.  Thus  in  this 
process,  which  is  used  in  Norway  and  is  known  as  the  Birkelund 
and  Eyde  process,  nitric  oxide  NO  is  formed  at  the  very  high 
temperature  of  the  arc.  It  is  the  high  temperature  secured  by 
means  of  the  electric  arc,  and  not  an  electrical  effect,  that 
causes  the  oxygen  and  nitrogen  to  unite.  When  the  nitric 
oxide  is  then  treated  with  air  and  water  in  a  tower  filled  with 
moist  coke,  nitrogen  dioxide  and  nitric  acid  form,  thus :  — 

2  NO  +  02  ;£  2  N02. 
H20  +  3  N02  5±  NO  +  2  HNOS. 

All  of  the  reactions  involved  in  the  process  are  reversible,  so 
that  in  order  to  have  them  go  to  completion  from  left  to  right 
as  far  as  possible,  the  products  formed  are  rapidly  removed  by 
condensation  and  solution.  The  dilute  nitric  acid  thus  obtained 
is  neutralized  with  lime,  and  the  calcium  nitrate  formed  is  sold 
as  a  fertilizer.  This  process  of  making  nitric  acid  is  of  special 
importance,  because  the  large  demands  made  upon  the  deposits 


164  OUTLINES   OF   CHEMISTRY 

of  Chili  saltpeter  annually  will  ere  long  exhaust  this  source  of 
supply,  though  new  deposits  have  been  found  of  recent  years  in 
the  same  region.  About  one  and  a  half  million  tons  of  Chili 
saltpeter  have  been  used  annually  in  recent  years.  The  salt  is 
used  as  a  fertilizer  to  a  large  extent,  but  it  is  also  employed  in 
making  nitric  acid,  which  is  used  in  the  manufacture  of  explo- 
sives, dyestuffs,  medicinal  chemicals,  nitrates  of  metals,  etc. 
Upwards  of  100,000  tons  of  nitric  acid  are  used  annually  in  the 
chemical  industries  of  the  world. 

Properties  of  Nitric  Acid.  —  Pure  nitric  acid  is  a  colorless 
liquid  which  boils  at  86°  with  partial  decomposition,  as  stated 
above.  It  is  a  monobasic  acid  whose  composition  is  expressed 
by  the  formula  HNO3.  On  distilling  the  acid  under  diminished 
pressure,  this  decomposition  may  be  avoided,  and  this  is  actually 
done  in  the  manufacture  of  pure  nitric  acid.  On  cooling,  nitric 
acid  forms  colorless  crystals  that  melt  at  42°.  The  acid  that  is 
sold  in  the  market  as  concentrated  nitric  acid  is  a  68  per  cent 
solution.  It  has  a  constant  boiling  point,  which  is  120.5°,  and 
a  specific  gravity  of  1.414  at  15°.  The  composition  of  this  con- 
stant boiling  solution  changes  when  it  is  distilled  under  dimin- 
ished pressure  (compare  hydrochloric  acid) ;  the  solution  is 
consequently  not  regarded  as  a  chemical  compound. 

Nitric  acid  is  a  powerful  acid  which  fumes  in  the  air.  In 
aqueous  solutions  it  is  much  more  stable  than  when  pure.  The 
concentrated  acid  is  rather  an  unstable  substance.  It  is  slowly 
decomposed  in  sunlight  to  a  slight  extent,  the  yellow  color  de- 
veloped being  due  to  the  formation  of  nitrogen  dioxide,  which 
remains  in  solution.  At  about  280°  nitric  acid  decomposes, 
practically  completely,  into  nitrogen  dioxide,  water,  and  oxy- 
gen. Concentrated  nitric  acid  has  a  very  corrosive  action  on 
the  skin,  producing  painful  wounds  that  are  slow  to  heal. 
More  dilute  solutions  color  the  skin  yellow,  due  to  the  forma- 
tion of  iiitro  products.  The  effect  upon  wool,  linen,  silk,  and 
other  organic  substances  is  similar. 

When  nitric  acid  is  neutralized  with  bases,  nitrates  are 
formed.  These  salts  are  all  readily  soluble  in  water.  Nitric 
acid  does  not  attack  gold  or  platinum.  When  it  attacks  other  met- 
als, they  are  either  oxidized,  as  is  the  case  with  tin,  or  converted 
into  nitrates,  as  is  more  frequently  the  case.  So  zinc,  copper, 
iron,  magnesium,  when  treated  with  nitric  acid,  are  converted 


OXY-ACIDS   AND   OXIDES   OF  NITROGEN  165 

into  the  corresponding  nitrates.  There  is,  however,  no  concom- 
itant evolution  of  hydrogen,  as  when  these  metals  are  attacked 
with  hydrochloric  acid,  for  instance ;  for  the  hydrogen  at  once 
attacks  the  nitric  acid,  reducing  it  commonly  to  nitric  oxide 
NO  and  water.  With  metals  like  zinc,  iron,  and  magnesium, 
the  temperature  and  concentration  of  the  acid  and  resulting 
solutions  may  be  regulated  so  as  to  secure  a  very  gradual  reduc- 
tion of  the  nitric  acid,  the  products  being  successively  nitrogen 
dioxide  NO2,  nitrous  acid  HNO2,  nitric  oxide  NO,  nitrous  oxide 
N2O,  nitrogen  N2,  hydroxylainine  NH2OH,  and  ammonia  NH3. 
As  nitrogen  dioxide  and  nitrous  acids  are  readily  reduced  to 
nitric  oxide,  the  latter  is  generally  formed  when  nitric  acid  acts 
on  a  metal,  thus  :  — 

3  Zn  +  8  HNO3  =  3  Zn(NO3)2  +  4  H2O  +  2  NO. 

Nitric  acid  is  a  powerful  oxidizing  agent  and  will  convert  many 
of  the  non-metals  into  their  highest  oxidation  products  with  ease. 
Thus,  when  heated  with  nitric  acid,  phosphorus  is  oxidized  to 
phosphoric  acid,  sulphur  to  sulphuric  acid,  carbon  to  carbon 
dioxide.  A  glowing  stick  of  charcoal  thrust  into  concentrated 
nitric  acid  continues  to  burn  brightly. 

While  neither  nitric  nor  hydrochloric  acid  alone  attacks  gold 
or  platinum,  these  metals  are  readily  dissolved  in  a  mixture 
of  nitric  and  hydrochloric  acids.  This  mixture,  since  it  dis- 
solves gold,  the  "  king  of  metals,"  is  called  aqua  regia.  The 
action  depends  upon  the  fact  that  nitric  acid  oxidizes  hydro- 
chloric acid,  one  of  the  products  formed  being  chlorine,  which 
attacks  gold.  Aqua  regia  was  known  even  in  the  days  of 
alchemy,  for  Geber  dissolved  gold  in  a  solution  of  ammonium 
chloride  in  nitric  acid.  The  action  of  concentrated  nitric  and 
hydrochloric  acids  on  each  other  may  be  represented  thus :  — 
3  HC1  +  HN03  =  2  H20  +  NOC1  +  C12. 

The  compound  NOC1  is  called  nitrosyl  chloride.  It  occurs 
here  as  one  of  the  products  of  the  reaction,  but  it  does  not 
attack  gold. 

Nitrogen  Pentoxide.  —  When  nitric  acid  is  treated  with  phos- 
phorus pentoxide,  nitrogen  pentoxide  or  nitric  acid  anhydride 
is  formed.  The  action  consists  of  the  subtraction  of  water  from 
nitric  acid :  — 

2  HN03  +  P206  =  2  HP08  +  N206. 


166  OUTLINES   OF  CHEMISTRY 

In  this  process,  pure  nitric  acid  is  carefully  mixed  with  about 
an  equal  weight  of  phosphorus  pentoxide  in  the  cold,  and  the 
sirupy  mass  obtained  is  carefully  distilled.  Nitric  anhydride 
may  also  be  formed  from  silver  nitrate  and  chlorine,  thus :  — 

4  AgN08  +  2  C12  =  4  AgCl  +  2  N2O6  +  Or 

It  was  by  this  method  that  the  substance  was  discovered  by 
Deville  in  1849.  Nitrogen  pentoxide  consists  of  colorless  pris- 
matic crystals  that  melt  at  30°,  forming  a  dark  yellow  liquid. 
The  latter  boils  at  50°  with  concomitant  partial  decomposition. 
It  is  very  unstable,  readily  giving  off  a  portion  of  its  oxygen, 
thus : — 

2N206  =  4N02+02. 

The  decomposition  goes  on  slowly,  though  spontaneously,  at 
ordinary  temperatures.  When  rapidly  heated,  the  decomposi- 
tion proceeds  with  explosive  violence.  The  substance  cannot 
be  kept  long  in  any  case.  Dissolved  in  water,  nitric  anhydride 
N2O5  yields  nitric  acid. 

Nitric  Oxide.  —  Nitric  oxide  NO,  discovered  by  Priestley  in 
1772,  is  formed  by  the  action  of  copper,  silver,  mercury,  and 
many  other  metals  upon  a  solution  of  nitric  acid  of  about  30  to 
35  per  cent :  — 

3  Cu  +  8  HNO3  =  3  Cu(NO3)2  +  4  H2O  +  2  NO. 

The  temperature  should  be  kept  low  during  the  reaction,  as 
otherwise  nitrous  oxide  N2O  and  nitrogen  are  apt  to  form. 
Nitric  oxide  is  also  conveniently  produced  by  the  action  of  fer- 
rous chloride  or  sulphate  on  nitric  acid  in  presence  of  hydro- 
chloric or  sulphuric  acid,  thus  :  — 

HNO3  +  3  FeCl2  +  3  HC1  =  2  H2O  +  3  FeCl8  +  NO. 
2  HNO3  +  6  FeSO4  +  3  H2SO4  =  4  H2O  +  3  Fe2(SO4)3+2  NO. 

The  gas  is  colorless,  but  on  coming  in  contact  with  the  oxygen 
of  the  air  it  immediately  turns  brown,  due  to  the  formation  of 
nitrogen  dioxide :  — 

2NO  +  O2=2NO2. 

It  is  consequently  necessary  to  expel  the  air  from  the  apparatus 
before  collecting  the  gas,  which  may  be  done  over  water  since 
the  latter  dissolves  nitric  oxide  but  slightly. 

Nitric  oxide  is  a  colorless,  neutral  gas  which  is  1.039  times 


OXY-ACIDS   AND   OXIDES   OF  NITROGEN  167 

as  heavy  as  air.  Its  critical  temperature  is  —  94°,  and  its  criti- 
cal pressure  71.2  atmospheres.  Under  atmospheric  pressure  the 
liquid,  which  is  colorless,  boils  at  — 150°.  When  solidified, 
nitric  oxide  forms  colorless  crystals  that  melt  at  —  167°.  At  0° 
one  volume  of  water  absorbs  0.075  volume  of  the  gas,  and  at 
20°,  0.05  volume. 

Nitric  oxide  is  the  most  stable  of  the  oxides  of  nitrogen.  A 
lighted  candle  or  burning  sulphur  will  be  extinguished  when 
introduced  into  the  gas.  On 
the  other  hand,  burning  mag- 
nesium or  phosphorus  will 
continue  to  burn  in  the  gas 
with  great  brilliancy. 

On  heating  metallic  sodium 

. .     .  .  ^       .-n,.  FlG.  41. 

or  iron  in  nitric  oxide  (Fig. 

41),  these  metals  are  oxidized,  and  the  nitrogen  which  remains 

occupies  just  one  half  of  the  volume  of  the  nitric  oxide,  thus  :  — 

4Na  +    2  NO   =   2  Na2O   +     N2. 

(solid)  (2  volumes)  (solid)  (1  volume) 

3Fe    +    4  NO   =   Fe3O4     +      2  N2. 

(solid)  (4  volumes)  (solid)  (2  volumes) 

Knowing  the  specific  gravities  of  nitric  oxide  and  nitrogen,  and 
the  fact  that  2  volumes  of  nitric  acid  yield  1  volume  of  nitrogen, 
it  follows  that  in  nitric  oxide  14  grams  of  nitrogen  are  combined 
with  every  16  grams  of  oxygen.  As  nitric  oxide  is  15  times 
heavier  than  hydrogen,  its  molecular  weight  is  30.  The  for- 
mula of  nitric  oxide  is  therefore  NO. 

Nitric  oxide  may  be  used  to  detect  the  presence  of  free  oxy- 
gen in  a  mixture  of  gases  on  account  of  its  ability  to  form 
brown  fumes  NO2  with  oxygen.  In  solutions  of  ferrous  salts 
nitric  oxide  dissolves  readily,  forming  a  dark  brown  liquid. 
From  these  solutions  the  gas  is  expelled  by  heating.  It  is 
probable  that  the  solutions  contain  the  unstable  compound 
FeSO4  •  NO.  This  reaction  is  a  delicate  test  for  nitrates,  for,  as 
we  have  seen,  ferrous  salts  in  presence  of  free  acid  readily 
reduce  nitrates  to  NO,  which  then  gives  the  brown  color  with 
the  excess  of  the  ferrous  salt. 

Nitrogen  Dioxide  and  Tetroxide.  —  Nitrogen  dioxide  is  pro- 
duced by  heating  nitrates  of  the  heavy  metals  :  — 


168  OUTLINES  OF  CHEMISTRY 

2  Pb(N03)2  =  2  PbO  +  02  +  4  N02. 
2  Cu(N03)2  =  2  CuO  +  02  +  4  NO2. 

When    oxygen^   acts    on    nitric    oxide,   nitrogen    dioxide    is 
formed  :  — 

2  NO  +  02  =  2  N02. 

When  the  electric  spark  is  passed  through  a  mixture  of  oxygen 
and  nitrogen,  nitrogen  dioxide  forms:  — 


As  stated  under  nitric  acid,  this  reaction  takes  place  slowly  and 
is  ordinarily  very  incomplete.  Concentrated  nitric  acid  oxi- 
dizes nitric  oxide  to  nitrogen  dioxide,  consequently  the  latter 
is  formed  when  metals  like  copper  or  tin  are  acted  upon  by 
strong  nitric  acid  even  out  of  contact  with  the  air. 

At  ordinary  temperatures,  nitrogen  dioxide  is  a  gas  of  a 
dark  reddish  brown  color.  When  chilled  with  a  freezing 
mixture,  consisting  of  ice  and  common  salt,  the  dark  brown 
nitrogen  dioxide  becomes  much  lighter  in  color  and  condenses 
to  a  pale  yellow  liquid.  At  —30°  this  liquid  congeals,  yielding 
colorless  crystals  that  melt  at  —10°,  thus  forming  a  colorless 
liquid  which  is  fairly  stable  even  at  0°.  On  gently  warming 
this  liquid,  it  assumes  a  greenish  yellow  hue.  At  about  10°  it 
is  yellow  in  color,  at  15°  it  is  orange  colored,  and  at  higher 
temperatures  it  becomes  still  darker,  till  at  26°,  its  boiling  point, 
the  color  becomes  a  dark  reddish  brown.  On  lowering  the 
temperature,  these  changes  occur  in  the  reverse  order.  At  2° 
the  vapor  is  38  times  as  heavy  as  hydrogen,  while  at  140°  the 
vapor  is  only  23  times  as  heavy  as  hydrogen.  At  26°,  therefore, 
the  molecular  weight  would  be  76,  and  at  140°,  46.  Now  the 
formula  NO2  corresponds  to  a  molecular  weight  of  46,  con- 
sequently at  140°  the  gas  is  NO2.  But  the  double  formula 
N2O4  corresponds  to  a  molecular  weight  of  92,  so  that  at  26° 
the  gas  has  more  nearly  the  formula  N2O4.  The  vapor  density 
decreases  gradually  as  the  temperature  is  raised,  and  all  these 
facts  are  best  explained  by  assuming  that  at  low  temperatures 
the  molecules  are  N2O4,  which  are  colorless,  and  that  these 
decompose  gradually,  with  rise  of  temperature,  into  brown 


molecules  of  NO2,  thus:  — 


N204;£2N02. 

(1  vol.)  (2  vols.) 


OXY-ACIDS  AND   OXIDES   OF   NITROGEN  169 

At  26°,  therefore,  the  dissociation  will  have  progressed  to  the 
extent  of  about  34  per  cent,  as  may  be  computed  from  the 
densities  above  given  ;  while  at  140°  the  dissociation  is  practi- 
cally complete. 

When  nitrogen  dioxide  acts  on  water  in  the  cold,  nitrous 
and  nitric  acids  result,  as  already  mentioned  in  connection  with 
nitric  acid. 

On  passing  nitrogen  dioxide  through  a  red-hot  tube,  it  is 
decomposed  into  oxygen  and  nitric  oxide.  The  action  is 
reversible',  thus  :  — 

2  N02  ;£  2  NO  +  02. 

Nitrogen  dioxide  is  a  poisonous  gas  having  a  corrosive  action 
on  the  mucous  membranes.  It  is  also  a  strong  oxidizing  agent, 
and  consequently  will  support  the  combustion  of  many  sub- 
stances. 

Nitrous  Acid.  —  When  potassium  nitrate  is  heated,  it  loses  a 
portion  of  its  oxygen  and  is  converted  into  potassium  nitrite, 

thus  :  — 

2  KNO3  =  2  KNO2  +  O2. 

The  nitrite  may  also  be  formed  by  heating  saltpeter  with  lead 
or  copper,  thus  :  — 

KN03  +  Pb  =  PbO  +  KNO2. 

KNO3  +  Cu  =  CuO  +  KNO2. 

Potassium  nitrite  KNO2  is  a  salt  of  nitrous  acid  HNO2.  The 
latter  has  never  been  prepared  except  in  dilute  solutions  at  low 
temperatures.  On  attempting  to  isolate  nitrous  acid  from  a 
nitrite  by  treating  with  sulphuric  acid,  the  following  reaction 
occurs  :  — 

2  KN02  +  H2S04  =  K2S04  +  H2O  +  NO  +  NO2. 

It  is  possible  that  at  first  HNO2  is  set  free,  which  undergoes 
decomposition  into  nitrogen  trioxide,  that  is,  nitrous  acid  an- 
hydride N2O3,  and  water,  thus  :  — 

2  HN02  =  H20  +  N203. 
The  latter  then  decomposes  into  nitric  oxide  NO  and  nitrogen 

Na03  =  NO  +  NO2. 


dioxide  NO2,  thus  :  — 


170  OUTLINES  OF   CHEMISTRY 

By  dissolving  nitrogen  trioxide  (see  below)  in  water  at  0°, 
a  blue  solution  is  obtained  which  is  commonly  regarded  as  a 
solution  of  nitrous  acid  HNO2.  This  solution  readily  evolves 
nitric  oxide  with  concomitant  formation  of  nitric  acid,  thus  :  — 

3  HN02  =  H20  +  2  NO  +  HNO3. 

Thus  it  is  evident  that  nitrous  acid  is  very  unstable.  Its  salts, 
however,  are  fairly  stable.  They  may  be  formed  by  neutraliz- 
ing aqueous  solutions  of  nitrous  acid  with  bases,  or  by  reduc- 
ing nitrates.  In  rain  water  and  frequently  in  contaminated 
drinking  water,  nitrites  are  present. 

Nitrous  acid  may  act  as  a  reducing  agent,  for  it  will  take  up 
oxygen  and  form  nitric  acid.  Thus  it  will  reduce  a  potassium 
permanganate  solution  as  follows  :  — 

2  KMn04  +  3  H2SO4  +  5  HNO2 

=  K2S04  +  2  MnS04  +  3  H2O  +  5  HNO8. 

On  the  other  hand,  toward  substances  that  will  take  up  oxygen, 
nitrous  acid  plays  the  rdle  of  an  oxidizing  agent.  So  with 
hydriodic  acid,  the  following  reaction  occurs  :  — 

2  HN02  +  2  HI  =  2  NO  +  2  H2O  +  Ia. 

Nitrous  acid  is  of  importance  in  the  study  of  carbon  compounds, 
and  in  the  preparation  of  aniline  dyes.  Nitrites  are  readily 
distinguished  from  nitrates,  for  nitrites  evolve  the  charac- 
teristic brown  nitrogen  dioxide  fumes  when  acidified  with 
sulphuric  acid.  Furthermore,  a  dilute  solution  of  a  nitrite 
acidulated  with  sulphuric  acid  will  turn  starch  potassium  iodide 
paper  blue  ;  compare  the  last  equation  above. 

Nitrogen  Trioxide.  —  Nitrogen  trioxide  or  nitrous  anhydride 
N2O3  readily  decomposes  into  NO  and  NO2.  When  equal 
volumes  of  the  latter  gases  are  mixed  and  cooled  to  —  21°  in  a 
tube,  nitrous  anhydride,  a  deep  blue  liquid,  is  formed.  It 
slowly  decomposes  even  at  —  21°,  but  at  its  boiling  point  the 
decomposition  is  more  rapid.  The  dissociation  may  be  indi- 
cated thus  :  — 


Hyponitrous  Acid.  —  By  reducing  sodium  or  potassium  nitrate 
or   nitrite  with  nascent   hydrogen   formed   by  the   action   of 


OXY-ACIDS   AND   OXIDES  OF  NITROGEN  171 

sodium  amalgam  on  the  aqueous  solution,  a  salt  of  hyponitrous 
acid  may  be  obtained  thus  :  — 

2  KNOa  +  4  H  =  K2N202  +  2  H2O. 

When  the  potassium  hyponitrite  is  treated  with  sulphuric  acid, 
the  hyponitrous  acid  liberated  is  decomposed  into  nitrous  oxide 
and  water,  thus  :  — 

K2N202  +  H2S04  =  K2S04  +  H2O  +  N2O. 

The  reaction  is  not  reversible,  and  so  hyponifrous  acid  can- 
not be  obtained  by  dissolving  N2O  in  water.  Free  hyponitrous 
acid  may  be  obtained  by  first  making  the  silver  salt  Ag2N2O2 
and  decomposing  this  by  means  of  hydrochloric  acid,  thus 
forming  insoluble  silver  chloride  AgCl  and  the  free  acid.  The 
latter  may  also  be  obtained  by  the  oxidation  of  hydroxylamine 
NH2OH  by  means  of  nitrous  acid,  thus :  — 

NH2OH  +  HN02  =  H2N2O2  +  H2O. 

The  anhydrous  acid  forms  transparent  crystalline  plates  which 
are  highly  explosive.  On  exploding,  the  acid  decomposes  into 
water,  nitrogen  and  oxygen ;  while  on  slow  decomposition  at 
room  temperatures  in  aqueous  solution  nitrous  oxide  and  water 
are  formed.  The  aqueous  solution,  however,  is  more  stable. 

Nitrous  Oxide.  —  On  heating  ammonium  nitrate,  nitrous 
oxide  N2O  and  water  are  produced,  thus :  — 

NH4NO3=2  H2O  +  N2O. 

A  mixture  of  ammonium  chloride  and  saltpeter  may  be  sub- 
stituted for  the  ammonium  nitrate,  for  thus  potassium  chloride 
and  ammonium  nitrate  are  formed,  and  the  latter  then  decom- 
poses on  heating  as  represented  above. 

Nitrous  oxide  is  a  colorless,  neutral  gas  which  is  1.52  times 
as  heavy  as  air.  It  has  a  sweetish  odor  and  taste,  and  when 
inhaled  it  produces  peculiar  symptoms  that  frequently  are 
accompanied  by  fits  of  hysterical  laughing ;  whence  its  name, 
laughing  gas.  On  continued  inhalation,  it  produces  insensi 
bility,  and  hence  the  gas  has  been  used  as  an  anaesthetic  in 
dental  operations.  The  gas  cannot  take  the  place  of  oxygen 
in  respiration,  however,  and  if  inhaled  for  a  long  time  death 
results. 


172  OUTLINES  OF   CHEMISTRY 

Nitrous  oxide  is  much  more  soluble  in  cold  than  in  warm 
water;  so  at  0°,  1.30  volumes  are  absorbed  by  1  volume  of 
water,  while  at  25°  only  0.59  volume  is  absorbed.  For  this 
reason  the  gas  is  collected  over  warm  water. 

Nitrous  oxide  boils  at  -  89.5°.     The  solid  melts  at  -  102.7°. 

It  may  be  obtained  in  the  market  in  compressed  form  in 
steel  cylinders.  A  glowing  splinter  burns  in  nitrous  oxide 
as  in  pure  oxygen.  Similarly  phosphorus  and  sulphur  burn 
in  nitrous  oxide  as  in  oxygen,  the  products  in  all  cases  being 
oxides  and  free  nitrogen.  Nitrous  oxide  is,  however,  readily 
distinguished  from  oxygen  by  the  fact  that  nitric  oxide  and 
oxygen  form  red  fumes  NO2,  which  does  not  take  place  when 
nitrous  oxide  and  nitric  oxide  are  mixed. 

When  mixed  with  an  equal  volume  of  hydrogen,  nitrous 
oxide  explodes  on  ignition  with  an  electric  spark,  and  the 
volume  of  nitrogen  formed  is  equal  to  that  of  the  original 
hydrogen,  thus :  — 

H2  +  N20  =  H20  +  N2. 

(1  volume)  (1  volume)       (liquid)     (1  volume) 

On  heating  nitrous  oxide  with  sodium  (Fig.  41),  the  following 
reaction  takes  place  :  — 

2  Na  +  N20  =  Na20  +  N2. 

(solid)        (1  volume)          (solid)       (1  volume) 

The  volume  of  nitrogen  formed  equals  that  of  the  nitrous  oxide. 
From  this  fact  and  the  specific  gravities  of  the  gases  involved,  it 
follows  that  the  formula  of  nitrous  oxide  is  N2O. 

While  nitrous  oxide  is  a  good  oxidizing  agent,  exhibiting 
many  of  the  properties  of  oxygen,  it  is  not  as  energetic  as  the 
latter.  So  metals  will  not  rust  in  contact  with  moist  N2O 
as  in  contact  with  moist  oxygen.  A  feebly  burning  piece  of 
sulphur  or  phosphorus  will  be  extinguished  in  nitrous  oxide, 
though  when  these  substances  are  burning  strongly,  they  con- 
tinue to  burn  brilliantly  in  the  gas. 

General  Considerations.  In  ammonia,  nitrogen  has  a  valence 
of  three,  thus  :  — 


OXY-ACIDS   AND   OXIDES  OF  NITROGEN  173 

In  ammonium  salts  it  has  a  valence  of  five,  thus  :  — 


C1 

In  hydroxylamine,   in   hydrazine,  in   hydrazoic   acid,   and  in 
nitrous  oxide,  nitrogen  is  trivalent,  thus  :  — 

,H  ft  .H        \=  ®        N  =  N 

N^H  ;        NN-N/    ;      XN/;       \o/  . 

XOH          W  XH 


(hydroxylamine)  (hydrazine)  (hydrazoic  acid)  (nitrous  oxide) 

At  first  sight  one  might  be  inclined  to  the  view  that  nitrogen 
is  univalent  in  N2O,  and  that  this  compound  is  analogous  to 
water,  the  two  hydrogen  atoms  of  which  are  replaced  by  nitro- 
gen atoms.  However,  the  ease  with  which  nitrous  oxide  parts 
with  its  oxygen  and  forms  free  nitrogen  speaks  for  the  above 
formula.  If  in  nitrous  oxide  the  oxygen  is  replaced  by  the 
bivalent  group  =  NH,  called  the  imide  group,  hydrazoic  acid 
results. 

In  the  salts  of  hydroxylamine  and  of  hydrazine,  nitrogen  is 
quinquivalent  as  in  the  ammonium  salts. 

In  nitrogen  pentoxide  and  nitric  acid  nitrogen  has  a  valence 
of  five,  thus  :  — 

^° 

N=0  ^O 

V),  and  N=O 


N)  -  H 


In  nitrogen  dioxide,  nitrogen  is  tetravalent,  thus  :  — 

O  =  N  =  O. 

In  nitrogen   tetroxide,  formed   by  chilling  nitrogen  dioxide, 
nitrogen  has  been  regarded  as  quinquivalent,  thus  :  — 


>o 

This  formula  is  not  generally  accepted,  however. 


174  OUTLINES  OF  CHEMISTRY 

In  nitrogen  trioxide  and  nitrous  acid  nitrogen  is  trivalent, 
thus :  — 

.O 

O    and  N— OH. 


In  nitric  oxide,  nitrogen  is  bivalent,  thus  :  — 


Finally  in  hyponitrous  acid  (N-O-H)2,  as  in  N2O,  nitrogen 
has  at  times  been  considered  as  univalent.  However,  the  com- 
pound NOH  is  not  known.  Attempts  to  isolate  it  have  always 
yielded  (NOH)2,  the  constitution  of  which  is  best  represented 
by  regarding  nitrogen  as  trivalent,  thus  :  — 

HO  -  N  =  N  -  OH. 

"From  this  compound,  water  readily  splits  off,  forming  nitrous 
oxide  thus  :  — 

N  =  N 

\o/'         :  ||| 

The  valence  of  nitrogen  thus  exhibits  a  relatively  great  range 
of  variation  in  different  compounds. 

We  may  regard  all  the  oxides  and  oxy-acids  of  nitrogen  as 
derived  from  hypothetical  hydroxides  by  successive  elimination  of 
water,  as  the  following  table  shows,  in  wjiich  all  the  compounds 
except  those  in  the  first  column  are  known  :  — 

2  N(OH)5  minus  4  H2O  =  2  HNO3  ;  2  HNO3  minus  H2O  =  N2O5. 
2N(OH)4minus4H2O  =  N2O4;        N2O4  yields  2  NO2. 
2  N(OH)3  minus  2  H2O  =  2  HNO2  ;  2  HNO2  minus  H2O  =  N2O3. 
2N(OH)2minus2H2O  =  2NO;        ......... 

2  NOH      .    -.     ,     .     .      N2O2H2;  N2O2H2  minus  H2O  =  N2O. 

There  is  also  a  striking  similarity  between  the  oxy-acids  of 
nitrogen  and  their  salts  on  the  one  hand,  and  the  oxy-acids  of 
the  halogens  and  their  salts  on  the  other  hand.  This  similarity, 
which  is  evident  from  the  following  table,  is  not  a  mere  simi- 
larity of  formulae,  for  the  compounds  themselves  exhibit  anal- 
ogies in  their  crystal  forms,  solubility  in  water,  and  general 


OXY-ACIDS  AND   OXIDES  OF  NITROGEN  175 

stability  on  heating  and  on  treatment  with  reagents.      Only 
known  compounds  are  included  in  the  table. 

N20  .  .  .  H2N202 .  .  .  Na2N2O2         C12O  .  .  .  HC1O  .  .  .  NaClO. 
NO 


N2O3 . . .  HNO2 NaNO2 NaClO2 . 

NO2     .     .:V    ;    ',  .  .     .     .          C102    . 

N205 . . .  HN03 NaNOg  I2O6  .  .  .  HC1O3  . . .  NaClO3 . 

C1207  ..HClO4...NaC104. 


CHAPTER  XIII 

SULPHUR,   SELENIUM,  AND  TELLURIUM 

Occurrence  and  Preparation  of  Sulphur.  —  Sulphur  has  been 
known  since  ancient  times,  for  it  occurs  in  nature  in  the  uneom- 
bined  state,  especially  in  the  vicinity  of  active  or  extinct  volca- 
noes. Thus  in  Italy,  Sicily,  Spain,  Poland,  Egypt,  Iceland, 
California,  the  Yellowstone  Park,  China,  and  India,  sulphur  is 
found  native.  As  a  result  of  volcanic  action  sulphur  probably 
is  formed  by  the  reduction  of  hydrogen  sulphide  H2S  by  sul- 
phur dioxide  SO2,  thus  :  — 

2H2S  +  S02  =  2H20  +  3S. 

Sulphur  also  occurs  in  sedimentary  deposits,  where  it  is  formed 
as  a  product  of  the  decay  of  certain  bacteria  and  algse  which 
are  able  to  store  up  this  substance  in  their  organisms  in  form 
of  minute  particles.  This  sulphur  originates  from  deposits  of 
gypsum,  from  which  it  is  liberated  as  hydrogen  sulphide  as  the 
result  of  cellulose  fermentation.  This  hydrogen  sulphide  is 
then  taken  up  by  the  algse  and  bacteria,  which  convert  it  into 
sulphates ;  but  in  this  process  they  store  up  a  reserve  stock 
of  free  sulphur  in  their  bodies.  The  sulphur  which  is  found 
in  sedimentary  deposits  then  really  occurs  from  the  oxidation 
of  hydrogen  sulphide  through  the  action  of  these  organisms, 
thus :  — 

H2S  +  0=H20  +  S. 

Some  of  the  sedimentary  sulphur  is,  however,  probably  also 
formed  by  direct  oxidation  of  hydrogen  sulphide  by  the  oxygen 
of  the  air.  Especially  rich  sedimentary  deposits  of  sulphur 
occur  in  Texas  and  Louisiana,  where  by  means  of  superheated 
steam  the  sulphur  in  the  lower  strata  is  melted  and  forced  up 
to  the  surface  in  the  liquid  state.  On  account  of  this  rich 
deposit  of  sulphur,  the  amount  produced  in  the  United  States 
in  1910  was  255,534  long  tons,  which  is  about  half  of  the 
world's  annual  production  of  sulphur. 

176 


SULPHUR,    SELENIUM,    AND   TELLU-RIUM 


177 


Sulphur  further  occurs  as  hydrogen  sulphide  in  the  waters  of 
sulphur  springs  and  in  the  air  near  active  volcanoes,  where  sul- 
phur dioxide  is  also  frequently  found.  In  combination  with 
metals,  sulphur  occurs  as  sulphides,  as  in  galenite  PbS,  pyrite 
FeS2,  zinc  blende  ZnS,  cinnabar  HgS,  and  copper  pyrite  CuFeSa. 


Fia.  42. 

It  is  also  found  in  form  of  sulphates  of  various  metals.  Thus  fer- 
rous sulphate  FeSO4,  lead  sulphate  PbSO4,  heavy  spar  BaSO4, 
are  found  in  nature  ;  but  above  all,  gypsum  CaSO4  •  2  H2O  and 
anhydrite  CaSO4  are  found  in  very  extensive  deposits.  The 
amount  of  gypsum  produced  in  the  United  States  alone  in  1910 
was  2,379,057  tons.  Sulphur  occurs  in  small  quantities  in  com- 


178  OUTLINES   OF  CHEMISTRY 

bination  with  other  elements  in  nearly  all  plant  and  animal 
tissues,  for  it  is  a  constituent  of  albumen.  So  it  is  found  par- 
ticularly in  muscles,  hair,  nails,  hoofs,  and  horns.  In  urine 
sulphur  is  found  as  sulphates.  In  some  plants,  like  mustard, 
onions,  garlic,  and  skunk  cabbages,  it  enters  into  odoriferous 
compounds  that  have  an  irritating  action  on  the  mucous  mem- 
branes and  the  skin. 

The  preparation  of  sulphur  from  the  native  deposits  consists 
of  melting  it  out  of  contact  of  the  air  and  thus  freeing  it  from 
the  gypsum,  calcium  carbonate,  sand,  etc.,  with  which  it  is  com- 
monly contaminated.  Thus,  a  raw  material  about  90  per  cent 
pure  is  obtained,  which  is  placed  in  cast-iron  retorts  and  distilled 
(Fig.  42).  The  vapors  enter  brick  chambers,  where  they  are 
condensed  on  the  cold  walls  in  form  of  fine  powder  which  is  placed 
on  the  market  as  flowers  of  sulphur.  As  the  walls  finally 
become  hot  the  sulphur  melts  and  collects  on  the  bottom  of  the 
chamber,  where  it  is  drawn  off  from  time  to  time  and  cast  into 
sticks  in  moist,  wooden,  slightly  conical  molds.  In  this  form  it 
is  called  roll  sulphur  or  brimstone.  Sulphur  is  also  prepared 
by  heating  pyrites  FeS2  and  condensing  the  product.  It  is 
further  prepared  from  the  waste  liquors  of  the  Le  Blanc  soda 
process  (which  see),  and  from  the  sulphide  of  iron  secured  as  a 
by-product  in  purifying  illuminating  gas. 

Properties  of  Sulphur. — Native  sulphur  and  roll  sulphur 
form  lemon-yellow  crystals,  of  specific  gravity  2.06,  belonging 
to  the  orthorhombic  system  (Fig.  43).  When 
heated,  this  rhombic  sulphur  melts  at  114.5° 
to  a  mobile,  light  yellow  liquid,  which  on 
further  heating  to  160°  becomes  dark  brown 
and  viscous.  In  the  neighborhood  of  200° 
the  viscosity  is  so  great  that  the  vessel  in 
which  the  sulphur  is  contained  may  be 
turned  bottom  upward  without  causing  the 
FlG-  43-  sulphur  to  run  out.  On  still  further  heat- 

ing, the  viscosity  of  the  liquid  diminishes,  but  its  color  remains 
dark  brown.  At  400°  the  liquid  is  quite  mobile,  and  at  450°  it 
boils,  emitting  a  heavy,  dark  brown  vapor. 

When  sulphur  is  melted  in  a  crucible  and  the  mass  is  allowed 
to  cool  till  a  crust  forms  over  the  top  of  the  liquid,  and  the 
latter  is  then  poured  out  through  a  hole  punctured  in  the  crust, 


SULPHUR,   SELENIUM,  AND  TELLURIUM  179 

it  is  found  that  the  walls  of  the  crucible  are  lined  with  needle- 
like,  almost  colorless  crystals  of  sulphur  that  belong  to  the 
monoclinic  system.  These  crystals  of  monoclinic  sulphur  melt 
at  119°.  They  have  a  specific  gravity  of  1.96.  On  standing 
they  very  slowly  change  to  crystals  of  the  orthorliombic  sys- 
tem. The  rhombic  crystals  are  thus  the  stable  ones  at  ordinary 
temperatures^  whereas  the  monoclinic  crystals  are  stable  at  high 
temperatures.  The  temperature  at  which  the  transition  from 
the  one  form  to  the  other  takes  place  is  96.5°;  at  this  point 
both  rhombic  and  monoclinic  sulphur  remain  side  by  side  in 
equilibrium  with  each  other  without  change.  Below  the  tran- 
sition point  all  passes  over  into  rhombic  sulphur,  while  slightly 
above  that  point  all  is  converted  into  the  monoclinic  variety. 

When  sulphur  heated  almost  to  the  boiling  point  is  poured 
into  cold  water,  an  elastic  mass  is  formed  which  is  called  plas- 
tic sulphur.  After  a  few  days  it  loses  its  plasticity  and  be- 
comes hard,  but  for  a  while  it  remains  non-crystalline,  that  is, 
amorphous.  This  amorphous  sulphur  is  practically  insoluble  in 
all  solvents  ;  however,  it  very  gradually  passes  over  into  rhom- 
bic sulphur.  This  is  soluble  in  carbon  disulphide  to  the  extent 
of  about  40  parts  in  100  at  room  temperature.  On  evapora- 
tion, it  may  again  be  obtained  from  this  solution  in  rhombic 
form.  Rhombic  sulphur  is  also  soluble  to  a  slight  extent  in 
liquids  like  alcohol,  ether,  turpentine,  fats,  and  linseed  oil. 
Flowers  of  sulphur  dissolve  only  partially  in  carbon  disulphide. 
They  are  a  mixture  of  amorphous  and  rhombic  sulphur. 

Substances  which  are  able  to  crystallize  in  two  different  systems 
are  called  dimorphous.  This  property  is  not  uncommon.  In 
passing  from  the  monoclinic  to  the  rhombic  form,  sulphur 
slowly  evolves  heat. 

Precipitated  sulphur,  or  milk  of  sulphur,  is  prepared  by  add- 
ing an  acid  to  a  polysulphide  like  K2S5 :  — 

K2S5  -f  2  HC1  =  2  KC1  -h  H2S  +  4  S. 

Thus  formed,  it  is  a  grayish  white  powder,  which  is  used  in 
medicine.  Precipitated  sulphur  is  soluble  in  carbon  disulphide. 
Sulphur  is  thus  a  polymorphous  substance.  The  ability  of 
an  element  to  occur  in  different  forms  has  been  called  allotro- 
pism,  and  so  the  different  forms  of  sulphur  are  sometimes  called 
the  allotropic  forms  of  sulphur.  Their  existence  has  been 


180  OUTLINES  OF   CHEMISTRY 

explained  by  assuming  that  the  molecules  of  the  different  modi- 
fications contain  a  different  number  of  atoms,  similar  to  the 
case  of  oxygen  and  ozone.  It  is  doubtful,  however,  whether 
the  cases  are  similar. 

Sulphur  is  insoluble  in  water  and  is  devoid  of  taste  and 
smell.  In  contact  with  moist  air  it  very  slowly  oxidizes  super- 
ficially and  passes  into  solution  as  sulphuric  acid.  Sulphur 
combines  with  many  metals  and  non-metals,  forming  sulphides. 
Heated  together  with  iron  or  copper,  for  instance,  the  union 
takes  place  with  evolution  of  light  and  heat.  In  the  air  and  in 
oxygen,  sulphur  burns  to  sulphur  dioxide  SO3,  which  in  pres- 
ence of  platinum,  black  will  take  on  more  oxygen  and  form  SO3. 
The  atomic  weight  of  sulphur  is  32.07.  Investigations  of  the 
vapor  density  of  sulphur  show  that  at  diminished  pressure  and 
low  temperatures  the  molecular  formula  of  sulphur  is  S8, 
whereas  at  800°  to  1000°  the  density  corresponds  to  the  formula 
S2.  There  is  a  gradual  decomposition  of  the  molecules  from 
S8  to  S2  as  the  temperature  rises.  In  the  neighborhood  of 
2000°  the  S2  molecules  are  further  largely  dissociated  into 
monatomic  molecules  S. 

Uses  of  Sulphur.  —  Sulphur  is  used  in  the  manufacture  of 
sulphuric  acid  and  sulphur  dioxide,  the  latter  being  used  as  a 
bleaching  and  disinfecting  agent.  Sulphur  is  also  used  in 
making  black  gunpowder,  fireworks,  vulcanized  caoutchouc, 
and  hard  rubber.  In  medicine  it  is  employed  as  a  specific. 

Crystals  and  Crystal  Systems.  —  Many  substances  are  able 
to  assume  the  crystalline  state.  Crystals  are  generally  formed 
by  allowing  liquids  to  congeal  or  solutions  to  evaporate  to  a 
point  at  which  the  dissolved  substances  separate  out.  Crystals 
may,  however,  also  be  formed  when  vapors  condense,  as  in  the 
sublimation  of  iodine  or  sulphur ;  or  they  may  form  gradually 
from  amorphous,  solid  substances,  as  in  the  case  of  the  conver- 
sion of  amorphous  sulphur  to  rhombic  sulphur.  There  are 
many  substances  which,  like  sulphur,  are  known  in  both  the 
crystalline  and  amorphous  states;  others  have  never  been 
found  in  crystalline  condition,  like  cellulose  and  dextrine ; 
whereas  still  others,  like  water,  are  always  crystalline  when 
solid.  Crystalline  substances  are  said  to  have  crystallizing 
power,  whereas  those  substances  that  are  only  known  in  amor- 
phous form  are  said  to  be  devoid  of  crystallizing  power.  We 


SULPHUR,    SELENIUM,   AND   TELLURIUM  181 


FIG.  44. 


FIG.  45. 


FIG.  46. 


FIG.  47. 


FIG.  48. 


FIG.  49. 


FIG.  50. 


FIG.  51. 


FIG.  52. 


^x  ^ 

PV 
^ 

f\          / 

> 

t            i 

\x               / 

NJl                       V 

FIG.  53. 

FIG.  54. 


182 


OUTLINES   OF  CHEMISTRY 


do  not  know  of  what  this  tendency  to  form  crystals  really  con- 
sists, much  less  are  we  able  to  measure  or  compare  quantita- 
tively the  crystallizing  power  of  various  substances. 

The  most  striking  external  characteristic  of  a  crystal  is  its 
regularity  of  form.  A  study  of  crystals  has  led  to  the  conclu- 
sion that  a  crystal  is  a  solid  bounded  by  plane  faces  which  are  the 
outcome  of  a  regular  internal  arrangement  of  the  molecules.  So 


FIG.  55. 


FIG.  56. 


FIG.  57. 


FIG.  58. 


FIG.  59. 


FIG.  GO. 


the  hardness,  color,  index  of  refraction,  crushing  strength,  resist- 
ance to  corrosion  by  chemical  agents,  etc.,  may  vary  as  different 
directions  in  one  and  the  same  crystal  are  considered. 

It  has  been  found  that  all  known  crystals  may  be  classified  into 
six  crystal  systems,  according  to  their  symmetry. 

All  crystals  whose  faces  may  be  referred  to  a  system  of  three 
axes  of  equal  length  and  at  right  angles  to  one  another  are 
said  to  belong  to  the  isometric  or  regular  system.  Some  com- 
mon forms  are  shown  in  Figs.  44  to  54.  These  crystals  may 
have  nine  so-called  planes  of  symmetry,  a  plane  of  symmetry 
being  a  plane  which  cuts  a  crystal  in  two  halves  that  are  to 
each  other  as  an  object  is  to  its  reflection  in  a  mirror.  Many 


SULPHUR,    SELENIUM,    AND   TELLURIUM 


substances  crystallize  in  the  regular  system.  Among  these  are 
common  salt,  alum,  fluorspar,  galena,  pyrite,  garnet,  diamond, 
gold,  silver,  mercury,  and  copper. 

Crystals  whose  planes  may  be  referred  to  a  system  of  three 
axes,  of  which  but  two  are  of  equal  length  but  all  at  right 
angles  to  one  another,  are  said  to  belong  to  the  tetragonal  or 
quadratic  system,  in  which  there  are  five  planes  of  symmetry 
possible.  Figures  55  to  60  show  some  common  forms  of  crys- 


FIG.  64. 


FIG.  65. 


FIG.  66. 


FIG.  67. 


tals  of  this  system  as  they  occur  in  rutile,  titanium  dioxide 
TiO2  ;  in  cassiterite,  stannic  oxide  SnO2,  the  most  important 
ore  of  tin ;  and  in  calomel  HgCl. 

In  the  hexagonal  system,  the  forms  may  be  referred  to  four 
axes,  three  of  which  are  of  equal  length,  lie  in  the  same  hori- 
zontal plane,  and  bisect  one  another  in  a  point  so  as  to  form  six 
angles  of  sixty  degrees  each.  The  fourth  axis  is  either  longer 
or  shorter  than  the  others,  and  runs  through  their  point  of 
intersection  at  right  angles  to  the  horizontal  plane,  which 
bisects  the  vertical  axis.  In  this  system  there  are  seven  pos- 
sible planes  of  symmetry.  Figures  61  to  67  show  some  typical 


184 


OUTLINES   OF   CHEMISTRY 


crystals  of  the  hexagonal  system.  To  it  belong  the  crystal 
forms  assumed  by  many  important  substances,  like  water  H2O, 
quartz  SiO2,  calcium  carbonate  CaCO3,  Chili  saltpeter  NaNO3, 
and  calcium  phosphate  Ca3(PO4)2.  The  so-called  rhombohedral 


FIG. 


FIG.  69. 


FIG.  70. 


FIG.  71. 


division  of  the  hexagonal  system  in  particular  has  many  repre- 
sentatives. It  has  sometimes  been  termed  a  separate  system, 
the  trigonal  system. 

Crystal  forms  that  can  be  referred  to  a  system  of  three  axes, 
all  of  which  are  at  right  angles  to  one  another  but  of  unequal 
lengths,  are  said  to  belong  to  the  orthorhombic  or  rhombic  sys- 
tem, in  which  there  are  but  three  possible  planes  of  symmetry. 


\/ 


FIG.  72. 


FIG.  73. 


FIG.  74. 


Figures  68  to  71  exhibit  some  typical  rhombic  forms.  In  this 
system  crystallize  many  substances,  like  sulphur,  iodine,  olivine 
Mg2SiO4,  saltpeter  KNO3,  heavy  spar  BaSO4,  and  magnesium 
sulphate  MgSO4  •  7  H2O. 

In  the  monoclinic  or  monosymmetric  system  the  forms  are 
referred  to  a  system  of  three  axes  all  of  which  are  of  unequal 
length.  The  two  axes  that  lie  in  the  vertical  plane  bisect  each 
other  at  right  angles,  and  the  third  axis  is  bisected  at  the  point 


SULPHUR,   SELENIUM,  AND  TELLURIUM 


185 


FIG.  75. 


FIG.  76. 


of  intersection  of  the  other  two,  but  it  does  not  make  a  right 
angle  with  the  plane  in  which  the  other  two  axes  lie.  The 
angle  which  it  makes  with  that  plane  varies  in  different 
crystals.  In  this  system  there  is  but  one  plane  of  symme- 
try. Figures  72  to  74  show  some  representative  monoclinic 
forms.  Many  compounds  crystallize  in  this  system,  among 
which  are  monoclinic  sulphur, 
gypsum,  feldspar,  cane  sugar, 
Glauber's  salt  Na2SO4  - 10  H2O, 
copperas  FeSO4  •  7  H2O,  and 
borax  Na2B4O7  •  10  H2O. 

Finally,  in  the  triclinic  or 
asymmetric  system  the  forms 
are  referred  to  three  unequal 
axes  bisecting  one  another  in  a 
point  at  angles  that  are  unlike 
and  not  right  angles.  In  this  system  there  is  no  symmetry 
whatever.  Figures  75  and  76  show  some  triclinic  forms. 
Copper  sulphate  CuSO4  •  5  H2O,  plagioclase  feldspar  NaAlSi3O8 
(albite),  and  CaAl2Si2O8  (anorthite)  crystallize  in  the  triclinic 
system. 

Under  the  same  conditions  a  chemical  substance  always  crystal- 
lizes in  the  same  system. 

Most  substances  crystallize  in  but  one  system.  However, 
we  have  seen  that  under  different  conditions  one  and  the  same 
substance  may  crystallize  in  two  different  systems.  This  prop- 
erty is  called  dimorphism.  Thus,  sulphur  may  form  rhombic  or 
monoclinic  crystals ;  calcium  carbonate  CaCO3  may  form  hex- 
agonal or  rhombic  crystals ;  iron  pyrites  may  form  isometric  or 
rhombic  crystals.  These  substances  are  consequently  dimor- 
phous. Substances  that  have  similar  chemical  composition 
generally  crystallize  in  the  same  system  and  exhibit  the  same 
forms.  This  is  the  law  of  isomorphism,  discovered  by  Eilhard 
Mitscherlich.  So,  for  instance,  the  carbonates  CaCO3,  FeCO3, 
MgCO3  are  rhombohedral ;  the  chlorides  NaCl,  KC1,  NH4C1  are 
isometric. 

Hydrogen  Sulphide.  —  This  is  by  far  the  most  important  com- 
pound which  sulphur  forms  with  hydrogen.  The  elements 
unite  directly  with  each  other  at  higher  temperatures,  forming 
the  compound  whose  composition  and  vapor  density  are  repre- 


186  OUTLINES   OF  CHEMISTRY 

sented  by  the  formula  H2S.  So  when  a  current  of  hydrogen 
is  passed  over  heated  sulphur  in  a  tube,  H2S  is  formed;  also 
when  certain  sulphides  are  similarly  heated  in  a  current  of 
hydrogen,  thus  :  — 


The  common  way  of  preparing  the  gas  consists  of  treating  fer- 
rous sulphide  FeS  (made  by  heating  sulphur  and  iron  together) 
with  either  dilute  sulphuric  or  hydrochloric  acid  ;  — 

FeS  +  H2SO4  =  FeSO4  +  H2S. 
FeS  +  2  HC1  =  FeCl2  +  H2S. 

Instead  of  ferrous  sulphide,  which  is  the  cheapest,  other  sul- 
phides might  be  employed.  The  gas  may  also  be  prepared  by 
reduction  of  sulphuric  or  sulphurous  acids  with  nascent  hydro- 
gen:— 

H2SO3  +  6  H  =  3  H2O  +  H2S. 

In  nature  hydrogen  sulphide  occurs  in  sulphur  springs,  vol- 
canic gases,  and  wherever  organic  matter  is  decomposing,  as  in 
sewer  gas,  in  the  intestinal  gases,  and  in  some  pathological 
cases  in  urine. 

Hydrogen  sulphide  is  a  colorless  gas  which  is  1.19  times  as 
heavy  as  air.  It  boils  at  —  62°  and  melts  at  —  86°.  The  gas 
has  a  very  disagreeable  odor,  being  that  of  rotten  eggs,  in 
which  it  is  contained.  Hydrogen  sulphide  is  a  very  poisonous 
gas  and  overcomes  persons  and  animals  suddenly,  in  which  respect 
it  resembles  hydrocyanic  acid.  Inhaled  in  small  amounts, 
hydrogen  sulphide  produces  headache  and  at  times  vomiting. 
The  gas  is  combustible,  burning  with  a  blue  flame  to  water 
and  sulphur  dioxide  :  — 

2  H2S  +  3  O2  =  2  H2O  +  2  SO2. 

In  an  insufficient  amount  of  oxygen,  the  products  are,  in  part, 
water  and  sulphur  :  — 


In  water,  hydrogen  sulphide  is  but  slightly  soluble,  about 
3  volumes  being  absorbed  by  1  volume  of  water  at  ordinary 
temperature  and  pressure.  On  boiling  this  solution,  all  the 
gas  escapes.  On  standing  exposed  to  the  air,  the  gas  in  the 


SULPHUR,    SELENIUM,   AND   TELLURIUM  187 

solution  is  gradually  oxidized  to  water  and  sulphur  which 
separates  out  in  the  form  of  a  precipitate. 

When  chlorine,  bromine,  or  iodine  act  on  hydrogen  sulphide, 
the  latter  is  decomposed,  sulphur  being  liberated  and  hydro- 
halogen  being  formed,  so  for  instance :  — 

H2S  +  I2  =  2  HI  +  S. 

The  aqueous  solution  of  hydrogen  sulphide  is  feebly  acid 
toward  litmus,  and  in  many  ways  it  deports  itself  like  a  weak 
acid.  So  it  will  react  with  metals  even  at  room  temperature, 
forming  sulphides  and  hydrogen,  thus  :  — 

2  Ag  +  H2S  =  Ag2S  +  H2. 
Pb  +  H2S  =  PbS  +  H2. 

Furthermore  it  reacts  with  many  basic  oxides  and  hydroxides, 
thus : — 

PbO  +  H2S  =  PbS  +  H20. 

2  NH4OH  +  H2S  =  (NH4)2S  +  2  H2O. 

KOH  +  H2S==KSH+H20. 
2  KOH  +  H2S  =  K2S  +  2  H2O. 

The  sulphides  of  sodium  and  potassium  show  a  strong  alkaline 
reaction  toward  indicators.  They  are  salts  of  a  very  weak 
acid  with  a  strong  base,  and  hence  are  decomposed  by  water 
by  hydrolysis.  The  reaction,  which  is  reversible,  may  be 
written  thus :  — 

K2S  +  H20:£KSH  +  KOH. 

When  passed  through  a  red-hot  tube,  hydrogen  sulphide  is 
decomposed  to  hydrogen  and  sulphur.  It  thus  parts  readily 
with  its  hydrogen,  and  is  consequently  a  good  reducing  agent, 
as  is  evident,  for  instance,  from  the  fact  that  it  will  reduce 
sulphuric  or  nitric  acid,  thus  :  — 

H2SO4  +  H2S  =  2  H2O  +  SO2  +  S. 
2  HNO3  +  3  H2S  =  4  H2O  +  2  NO  +  3  S. 

Hydrogen  sulphide  is  a  very  important  reagent  in  chemical 
analysis,  for  while  the  sulphides  which  it  forms  with  metals 
like  sodium,  potassium,  calcium,  and  magnesium  are  soluble  in 
water,  other  sulphides  like  those  of  iron,  zinc,  and  nickel  are 
not  soluble  in  water,  but  soluble  in  dilute  acids,  and  still  other 
sulphides  like  those  of  arsenic,  copper,  and  lead  are  insoluble 


188  OUTLINES  OF   CHEMISTRY 

both  in  water  and  dilute  acids.  A  very  careful  study  of  these 
and  similar  properties  of  the  sulphides  of  the  metals  has  led  to 
a  system  by  means  of  which  the  metals  can  be  detected  and 
separated  when  they  occur  together. 

Polysulphides  and  Hydrogen  Persulphide. — When  sulphur 
is  added  to  a  solution  of  sulphide  of  potassium,  sodium,  calcium, 
ammonium,  etc.,  it  dissolves,  forming  polysulphides.  Thus, 
with  K2S  sulphur  may  form  compounds  varying  in  composition 
from  K2S  to  K2S5  according  to  the  amount  of  sulphur  dissolved. 
When  such  a  persulphide  is  gradually  added  to  a  very  dilute 
solution  of  hydrochloric  acid,  a  thick,  yellow  oil  of  disagree- 
able odor  separates  out  which  has  the  composition  H2S6,  no 
matter  what  the  sulphur  content  of  the  poly  sulphide  was, 
thus :  — 

2  K2S3  +  4  HC1  =  4  KC1  +  H2S  +  H2S5. 
4  Na2S2  +  8  HC1  =  8  NaCl  +  3  H2S  +  H2S6. 

Hydrogen  persulphide  bleaches  organic  dyestuffs.  It  reacts 
with  iodine,  forming  hydriodic  acid  and  sulphur.  It  gradually 
decomposes  into  hydrogen  sulphide  and  sulphur  on  standing. 

Comparison  of  Hydrogen  Sulphide  with  Water.  —  It  is  evident 
that  hydrogen  sulphide  and  water  possess  many  points  of 
analogy.  Thus  the  one  is  H-S-H  and  the  other  H-O-H.  With 
the  univalent  metals  they  form  hydrosulphides  MSH  and 
hydroxides  M OH,  respectively;  furthermore,  the  corresponding 
sulphides  M2S,  and  oxides  M2O,  are  also  known.  With  ele- 
ments of  higher  valence,  analogous  sulphides  and  oxides  are 
formed.  Thus  we  have  FeS  and  FeO,  P2O5  and  P2S5,  Sb2O3 
and  Sb2S3,  etc.  Again,  just  as  oxygen  and  hydrogen  form  a 
peroxide  H2O2,  so  sulphur  and  hydrogen  form  a  persulphide, 
which,  to  be  sure,  has  the  composition  H2S5.  We  shall  later 
see  further  points  of  resemblance  between  oxygen  and  sulphur 
in  their  chemical  behavior.  The  two  elements  indeed  belong 
to  the  same  family  group. 

Compounds  of  Sulphur  with  the  Halogens.  —  Fluorine  unites 
directly  with  sulphur  to  form  sulphur  hexafluoride  SF6,  which 
consists  of  white  crystals  that  melt  at  —  55°.  The  substance 
boils  but  slightly  above  its  melting  point.  The  gas  is  colorless, 
odorless,  tasteless,  and  practically  as  indifferent  toward  other 
reagents  as  nitrogen. 


SULPHUR,    SELENIUM,    AND   TELLURIUM  189 

When  dry  chlorine  is  passed  over  molten  sulphur  in  a  tubu- 
lated retort,  sulphur  monochloride  S2C12,  boiling  at  138°,  is 
formed.  It  is  a  fuming  yellowish  red  liquid  of  suffocating 
odor.  Its  specific  gravity  is  1.7.  It  dissolves  sulphur  readily  ? 
solutions  containing  over  60  per  cent  sulphur  being  obtainable. 
For  this  reason  sulphur  monochloride  is  used  in  preparing 
vulcanized  rubber.  Water  decomposes  sulphur  monochloride, 
thus :  — 

2  S2C12  +  2  H20  =  4  HC1  +  SO2  +  3  S. 

Sulphur  dichloride  SC12  is  formed  when  sulphur  monochlo- 
ride is  saturated  with  chlorine  in  the  cold.  It  is  an  oil  of  reddish 
brown  color  and  specific  gravity  1.6.  It  readily  decomposes  at 
64°,  yielding  sulphur  and  sulphur  monochloride.  It  is  also  de- 
composed by  water,  thus :  — 

2  SC12  +  2  H2O  =  4  HC1  +  SO2  +  S. 

Sulphur  tetrachloride  SC14  is  formed  by  saturating  sulphur 
dichloride  with  chlorine  at  temperatures  below  —  25°.  The 
substance  forms  crystals  which  melt  at  —  30°.  It  readily 
dissociates  above  —  22°,  the  decomposition  being  practically 
complete  at  +6°.  With  water  it  reacts  violently,  thus:  — 

SC14  +  2  H20  =  S02  +  4  HC1. 

With  bromine,  sulphur  forms  sulphur  monobromide  S2Br2,  a 
brownish  red  liquid  which  congeals  at  —  46°  and  boils  at  about 
200°,  accompanied  by  partial  decomposition. 

With  iodine,  sulphur  forms  sulphur  monoiodide  S2T2,  consist- 
ing of  dark  grayish  crystals  melting  at  60°,  and  also  sulphur 
hexaiodide  SI6,  which  forms  dark  crystals  that  readily  decom- 
pose on  standing,  yielding  free  iodine. 

Sulphur  Dioxide  and  Sulphurous  Acid.  —  When  sulphur  is 
burned  in  the  air  or  in  oxygen,  the  following  reaction  takes 
place :  — 

S  +  02  =  S02. 

The  resulting  sulphur  dioxide  occupies  the  same  volume  as  the 
oxygen,  which  may  be  demonstrated  by  means  of  the  apparatus 
of  Victor  Meyer  shown  in  Fig.  77.  The  sulphur  is  burned  in 
oxygen,  with  which  the  flask  has  been  filled.  On  cooling,  the 
manometer  indicates  that  the  volume  of  the  gas  in  the  appara- 
tus has  not  changed. 


190 


OUTLINES  OF   CHEMISTRY 


Sulphur  dioxide  is  a  colorless  gas  of  suffocating  odor.     It  is 
2.21  times  heavier  than  air.     It  may  readily  be  condensed  to  a 

liquid  at  ordinary  pressure  by 
cooling  to  —  10°.  Under  a  pres- 
sure of  about  two  atmospheres 
it  may  be  liquefied  at  room  tem- 
peratures. The  liquid  boils  at 

—  8°,    and   the   solid  melts    at 

—  76°.      Sulphur   dioxide  will 
not  support  combustion ;  never- 
theless, at  higher  temperatures 
many     metallic     oxides     unite 
vigorously  with  it  with  evolu- 
tion of  light,  thus:  — 

PbO2-j-SO2=PbSO4. 


FIG.  77. 


Besides  being  produced  by 
the  burning  of  sulphur,  sulphur 
dioxide  is  formed  by  heating  sulphides  of  certain  metals  in  the 
air ;  thus,  pyrite  acts  as  follows  :  — 

2  FeS2  +  11  O  =  Fe2O8  +  4  SO2. 

In  the  laboratory,  sulphur  dioxide  is  commonly  made  by 
heating  copper  turnings  with  concentrated  sulphuric  acid :  — 

2  H2S04  +  Cu  =  2  H20  +  CuS04  +  SO2. 

It  may  also  be  formed  by  heating  concentrated  sulphuric  acid 
with  carbon  or  sulphur  :  — 

2  H2SO4  +  C  =  2  H2O  +  CO2  +  2  SO2. 
2  H2SO4  +  S  =  2  H2O  +  3  SO2. 

When  dilute  sulphuric  acid  acts  on  sulphites,  sulphur  dioxide 
is  formed ;  also  when  metallic  oxides  are  heated  with  sulphur :  — 

NaHSO3  +  H2SO4  =  NaHSO4  +  SO2  +  H2O. 
2  MnO2  +  4  $  =  2  MnS  +  2  SO2. 
2  CuO  +  2  S  =  Cu2S  +  SO2. 

In  the  presence  of  water,  sulphur  dioxide  bleaches  many  organic 
coloring  matters.  Figure  78  shows  the  bleaching  of  flowers  by 
sulphur  dioxide  evolved  by  burning  sulphur.  This  bleaching 


SULPHUR,    SELENIUM,   AND   TELLURIUM 


191 


does  not  depend  upon  the  oxidation  of  the  dyes,  but  rather 
upon  their  union  with  the  sulphur  dioxide,  for  on  warming 
some  of  the  articles  thus  bleached  their  color  may  be  restored, 
In  other  cases,  the  bleaching  action  depends  upon  the  subtrac- 
tion of  oxygen  from  the  sub- 
stances. Sulphur  dioxide  is  used 
to  bleach  silk,  wool,  straw,  and 
other  fibers  that  would  be  de- 
stroyed by  means  of  chlorine.  It 
is  also  used  as  an  antiseptic  and 
disinfectant,  for  it  is  a  powerful 
germicide.  For  these  purposes  it 
may  now  be  obtained  in  liquid 
form  in  tin  cans. 

About  50  volumes  of  sulphur 
dioxide  are  dissolved  by  1  volume 
of  water  at  15°,  while  at  40°  but 
18.8  volumes  are  thus  absorbed. 
From  the  solution  all  of  the  sul-  FlG>  78' 

phur  dioxide  may  be  expelled  by  boiling.  The  solution  reacts 
acid  and  behaves  as  though  it  contained  sulphurous  acid  H2SO3, 
but  this  substance  has  never  been  isolated,  thus  :  — 


S0 


H20  =  H2S03. 


With  bases,  sulphurous  acid  forms  salts  called  sulphites,  thus  :  — 

H2SO3  +  NaOH  =  NaHSO3  +  H2O. 
HS0  +  2  NaOH  =  NaSO  +  2  HO. 


H2S03  +  Ca(OH)2  =  CaS03 


2  H2O. 


Sulphurous  acid  is  dibasic  in  character.  Both  the  acid  and 
the  normal  sulphites  of  the  alkali  metals  are  soluble  in  water, 
but  other  normal  sulphites  are  sparingly  soluble.  From  sul- 
phites, sulphur  dioxide  may  readily  be  regenerated  by  addition 
of  sulphuric  or  hydrochloric  acid.  This  fact  is  used  in  the 
detection  of  sulphites  in  chemical  analysis. 

Sulphur  dioxide  is  a  reducing  agent,  which  property  is  pos- 
sessed in  a  still  greater  degree  by  its  aqueous  solutions.  This 
is  because  sulphurous  acid  is  able  to  take  up  additional  oxy- 
gen readily,  thus  passing  over  into  sulphuric  acid.  Even  the 


192  OUTLINES   OF   CHEMISTRY 

oxygen  from  the  air  slowly  converts  sulphurous  acid  into  sul- 
phuric acid  in  solution,  thus  :  — 

2  H2S08  +  02  =  2  H2S04. 

Chlorine,  bromine,  or  iodine  rapidly  change  sulphurous  acid 
into  sulphuric  acid,  thus  :  — 

H2S03  +  H20  +  C12  =  H2S04  +  2  HC1. 
H2S03  +  H20  +  Ia  =  H2S04  +  2  HI. 

Sulphur  Sesquioxide.  —  This  compound  has  the  composition 
S2O3.  It  may  be  prepared  by  treating  molten  sulphur  trioxide 
SO3  with  pulverized  sulphur.  The  product  consists  of  bluish 
green  crystals.  With  fuming  sulphuric  acid  it  forms  a  blue 
solution.  Water  decomposes  the  sesquioxide  into  sulphuric 
acid  and  sulphur. 

Sulphur  Trioxide  and  the  Contact  Process  of  making  Sulphuric 
Acid.  —  Sulphur  trioxide  SO3  is  formed  by  heating  sulphates 
of  many  of  the  heavy  metals,  thus  :  — 


Oxygen  unites  but  very  slowly  with  SO2  to  form  SO3,  in  spite 
of  the  fact  that  the  union  is  accompanied  with  considerable 
evolution  of  heat.  However,  when  a  mixture  of  sulphur  diox- 
ide and  oxygen  is  passed  over  finely  divided  platinum,  the 
union  readily  takes  place,  the  action  being  practically  complete 
at  450°.  In  this  process,  the  platinum  remains  unchanged.  It 
acts  as  a  contact  or  catalytic  agent.  In  place  of  finely  divided 
platinum,  ferric  oxide  or  chromic  oxide  will  also  serve.  The 
residues  of  the  oxides  obtained  by  roasting  pyrites  are  some- 
times used  for  this  purpose.  The  sulphur  dioxide  obtained  by 
burning  sulphur  or  roasting  native  sulphides,  generally  pyrites, 
is  mixed  with  air  in  such  proportion  that  there  is  present  a 
large  excess  of  oxygen  beyond  what  is  needed  to  produce  sul- 
phur trioxide  according  to  the  equation  :  — 

2SO2+O2^±2SO3; 

for  this  reaction  is  a  reversible  one  and  the  presence  of  the 
excess  of  oxygen,  according  to  the  law  of  mass  action,  displaces 
the  equilibrium  toward  the  right.  The  temperature  should  be 
held  at  about  400°  to  450°,  for  at  higher  temperatures  the  sul- 
phur trioxide  dissociates,  that  is,  the  action  reverses.  The 


SULPHUR,    SELENIUM,   AND   TELLURIUM 


193 


gases  should  be  purified.  It  is  especially  necessary  that  they 
be  freed  from  dust  and  from  arsenic.  The  latter  is  generally 
present  in  the  gases  and  is  removed  by  means  of  steam.  Both 
the  residues  from  roasting  pyrites,  and  platinized  asbestus  are 
used  at  present  in  thus  preparing  sulphur  trioxide  by  what  is 
known  as  the  "contact  process."  The  bulk  of  this  sulphur 
trioxide  formed  is  used  in  making  sulphuric  acid,  and  to  this 
end  it  is  absorbed  in  sulphuric  acid  of  97  to  98  per  cent 
strength.  The  strength  of  the  acid  is  regulated  by  addition 
of  water.  Enormous  quantities  of  sulphuric  acid  are  now  pre- 
pared annually  by  the  contact  process,  both  in  Europe  and 


FIG.  79. 

America;  and  this  method,  the  success  of  which  on  a  commer- 
cial scale  is  due  to  the  labors  of  Knietsch  (1901),  has  to  a  large 
extent  displaced  the  lead  chamber  process  for  making  sulphuric 
acid,  at  least  so  far  as  making  concentrated  sulphuric  acid  is 
concerned.  On  a  small  scale,  in  the  laboratory,  sulphur  trioxide 
can  readily  be  made  by  means  of  the  apparatus  shown  in  Fig.  79. 
Sulphur  dioxide  from  a  generator  and  oxygen  from  a  tank 
are  passed  into  the  wash-bottle  It;  the  mixed  gases  then  pass 
through  the  drying  tube  T,  filled  with  pumice  soaked  in  sul- 
phuric acid,  and  finally  enter  the  tube  containing  the  asbestus, 
which  contains  finely  divided  platinum  heated  to  400°.  The 
SO3  formed  is  condensed  in  the  receiver. 

Sulphur  trioxide  is  also  formed  by  heating  fuming  sulphuric 


194  OUTLINES  OF   CHEMISTRY 

acid  or  warming  concentrated  sulphuric  acid  with  phosphorus 
pentoxide,  or  by  heating  sodium  or  potassium  pyrosulphate, 
thus : — 

H2S2O7=H2SO4+SO3. 
H2SO4  +  P2O6  =  SO3  +  2  HPO3. 

K2S2O7  =  K2SO4  +  SO3. 

Sulphur  trioxide  forms  long,  colorless,  prismatic  crystals  that 
melt  at  14.8°,  forming  a  colorless,  mobile  liquid  that  boils  at 
46°.  At  20°  the  specific  gravity  is  1.97.  Below  27°  sulphur 
trioxide  forms  sulphur  hexoxide  S2O6,  the  crystals  of  which 
look  like  long-fiber  asbestus  and  melt  at  50°.  On  further  heat- 
ing, it  passes  over  into  vapors  that  are  identical  with  those  of 
SO3,  i.e.  it  dissociates  into  SO3,  which  on  cooling  yields  a  liquid 
boiling  at  46°.  Sulphur  trioxide  has  great  affinity  for  water. 
It  fumes  strongly  in  the  air,  and  unites  with  water  with  great 
avidity  and  liberation  of  much  heat  which  forms  steam,  causing 
a  hissing  noise  as  the  substance  is  brought  into  contact  with 
water.  It  is  dangerous  to  bring  large  quantities  of  sulphur 
trioxide  into  contact  with  water  at  once,  for  the  heat  liberated 
causes  explosions.  At  temperatures  above  600°  sulphur  triox- 
ide dissociates  into  sulphur  dioxide  and  oxygen,  the  reaction 
being  practically  complete  at  1000°. 

Sulphuric  Acid  and  the  Lead  Chamber  Process.  —  Sulphuric 
acid  H2SO4  has  been  known  for  a  long  time.  The  alche- 
mists prepared  it  by  heating  ferrous  sulphate,  green  vitriol 
FeSO4  •  7  H2O,  hence  the  name  oil  of  vitriol.  This  process  was 
described  by  Basil  Valentine  in  1450,  who  also  prepared  the 
acid  bsy  burning  sulphur  in  presence  of  saltpeter.  In  1746 
Roebuck,  in  England,  made  use  of  the  principle  of  the  latter 
method  by  burning  sulphur  mixed  with  saltpeter  in  closed 
leaden  chambers  in  presence  of  moisture  which  absorbed  the 
gases,  forming  sulphuric  acid.  By  admitting  more  air  to  the 
chambers,  and  burning  more  sulphur  in  them,  additional  sul- 
phuric acid  was  formed,  and  so  on.  This  process  was  the 
beginning  of  what  is  to  the  present  day  known  as  the  lead 
chamber  process  of  the  manufacture  of  sulphuric  acid.  In  its 
essence  the  method  consists  of  oxidizing  sulphurous  acid 
H2SO3  to  sulphuric  acid  H2SO4,  by  means  of  nitric  a(  \d  and 
its  decomposition  products. 


SULPHUR,    SELENIUM,   AND   TELLURIUM  195 

In  practice,  the  manufacture  of  sulphuric  acid  loy  the  lead 
chamber  process  involves:  (1)  The  burning  of  sulphur  to  sul- 
phur dioxide,  either  by  using  sulphur  or  commonly  by  roasting 
native  sulphides  like  pyrite  FeS2,  copper  pyrite,  CuFeS2,  gale- 
nite  PbS,  zinc  blende  ZnS ;  (2)  the  oxidation  of  the  sulphur 
dioxide  in  presence  of  water  by  means  of  nitric  acid  and  nitro- 
gen dioxide-  one  of  its  decomposition  products ;  (3)  the  oxi- 
dation of  the  nitric  oxide  NO  formed  by  the  reduction  of  the 
nitric  acid  and  NO2 ;  and  (4)  the  concentration  of  the  sul- 
phuric acid  obtained.  In  the  roasting  of  the  native  sulphides 
mentioned,  the  latter  are  heated  in  a  current  of  air,  whereby 
sulphur  dioxide  and  the  oxides  of  the  metals  result.  The  nitric 
oxide  is  oxidized  to  NO2  by  means  of  oxygen  of  the  air.  We 
may  write  the  chemical  changes  involved  as  follows  :  — 

(1)  S  +  02=S02. 

(2)  3  S02  +  2  H20  +  2  HN03  =  3  H2SO4  +  2  NO. 

(3)  2  NO  +  H20  +  30  =  2  HNO3,  and 

(4)  NO  +  O  =  N02. 

(5)  S02  +  H20  +  N02  =  H2SO4  +  NO. 

Thus  it  will  be  seen  that  when  nitric  acid  acts  on  sulphur  diox- 
ide in  presence  of  moisture  (equation  2),  sulphuric  acid  and 
nitric  oxide  result.  The  latter  is  then  oxidized  by  oxygen 
from  the  air,  in  part  to  nitric  acid  (equation  3),  and  in  part 
to  nitrogen  dioxide  (equation  4).  The  nitric  acid  so  formed 
then  reacts  with  more  sulphur  dioxide,  according  to  equa- 
tion (2),  and  the  nitrogen  dioxide  oxidizes  sulphurous  acid 
according  to  equation  (5),  the  nitric  oxide  NO  formed  in  both 
cases  being  again  oxidized  by  oxygen,  and  in  turn  reduced  by 
sulphurous  acid  with  concomitant  formation  of  sulphuric  acid, 
and  so  on. 

While  the  above  equations  may  be  used  to  represent  what 
occurs  in  the  manufacture  of  sulphuric  acid,  the  actual  process 
is  no  doubt  of  more  complicated  character.  It  has  been  studied 
by  various  investigators,  among  whom  George  Lunge  holds  that 
a  compound  HO  -SO2-O(NO),  nitrosyl  sulphuric  acid,  is  formed 
in  the  chambers  during  the  process,  and  that  this  compound  is 
then  decomposed  by  water  with  resulting  formation  of  sulphuric 


196  OUTLINES   OF   CHEMISTRY 

acid  HO  •  SO2-  OH.    The  reactions  involved  in  this  explanation 
are  :  — 

(1)  SO2+  HN03=  HO  -  S02  -  0(NO). 

(2) 


The  nitrosyl  sulphuric  acid  is  then  again  decomposed  by  water, 
according  to  equation  (2),  and  so  on.  In  nitrosyl  sulphuric 
acid  we  have  the  univalent  —  N=O  group,  which  takes  the 
place  of  one  of  the  hydrogen  atoms  in  sulphuric  acid.  Now, 
in  the  ordinary  manufacture  of  sulphuric  acid,  when  things 
are  running  properly,  the  formation  of  nitrosyl  sulphuric  acid, 
which  consists  of  colorless  crystals  known  as  "chamber  crys 
tals,"  is  not  observed.  It  is  only  when  the  supply  of  water  is 
deficient  that  these  crystals  are  actually  formed,  for  they  are 
decomposed  by  water,  as  stated  above.  Although  there  is  dif- 
ference of  opinion  as  to  what  actually  occurs  in  the  details  of 
the  sulphuric  acid  manufacture,  the  changes  in  which  process 
are  undoubtedly  rather  complicated,  it  nevertheless  is  certain 
that  by  this  process  sulphurous  acid  is  completely  and  economi- 
cally converted  into  the  end  product,  sulphuric  acid.  The  oxides 
of  nitrogen  can  be  used  over  and  over  again,  though  of  course 
there  is  always  some  loss  of  the  latter  that  must  be  replenished. 
The  accompanying  Fig.  80  shows  in  diagrammatic  form  the 
essentials  of  a  lead  chamber  sulphuric  acid  factory.  In  the 
furnaces  F,  the  pyrites  and  other  native  sulphides  are  roasted  in 
a  current  of  air.  The  sulphur  dioxide  thus  produced  contains 
dust  carried  along  mechanically,  which  deposits  in  a  special 
long  dust  flue  in  which  the  gas  is  also  mixed  with  air  in  proper 
proportion.  The  gases,  which  are  at  a  temperature  of  about 
300°,  then  pass  into  the  Glover  tower  Cr.  This  is  a  structure 
about  10  meters  high  and  3  meters  in  diameter,  lined  inside  with 
sheet  lead  and  filled  with  acid  proof  stones,  over  which  dilute 
sulphuric  acid  containing  oxides  of  nitrogen  in  solution  contin- 
ually trickles  from  the  reservoir  on  top  of  the  tower.  This  acid 
is  derived  from  the  Gay-Lussac  tower  and  from  the  chambers, 
and  contains  also  some  nitric  acid,  which  has  been  added  to 
replace  the  oxides  of  nitrogen  that  are  inevitably  lost  during 
the  process  of  manufacture.  As  the  hot  gases  from  the  furnaces 
come  into  contact  with  this  sulphuric  acid  of  the  Glover  tower, 


SULPHUR,   SELENIUM,   AND  TELLURIUM 


197 


198  OUTLINES   OF  CHEMISTRY 

they  are  gradually  cooled  till  they  attain  a  temperature  of  about 
70°  when  they  reach  the  top.  At  the  same  time,  the  acid  is 
heated  up  and  thus  concentrated,  water  being  lost  which  is  car- 
ried off  with  the  gases  in  form  of  steam.  Again,  practically  all 
of  the  oxides  of  nitrogen  are  carried  off  by  the  gases,  which  when 
they  leave  the  tower  pass  into  the  first  lead  chamber  laden  with 
oxides  of  nitrogen  and  water  vapor.  The  acid  which  flows  from 
the  bottom  of  the  Glover  tower  contains  only  traces  of  oxides 
of  nitrogen  and  is  about  80  per  cent  strong.  There  are  com- 
monly three  lead  chambers,  so  connected  that  the  gases  enter 
at  the  top  of  each  and  pass  out  at  the  bottom.  In  these  cham- 
bers, which  often  have  a  volume  of  1000  cubic  meters  each,  the 
reactions  above  mentioned  take  place.  In  the  first  and  second 
chambers,  water  vapor  is  added  to  the  gases.  This  is  done 
either  by  blowing  in  steam  from  the  boiler,  or  by  forcing  water 
into  the  chambers  in  form  of  a  spray.  In  the  third  chamber  the 
gases  are  cooled,  and  they  then  pass  (charged  with  oxides  of 
nitrogen  regenerated  during  the  formation  of  sulphuric  acid  in 
the  chambers)  into  the  bottom  of  the  Gay-Lussac  tower.  This 
is  lined  with  lead  and  filled  with  coke  over  which  80  per  cent 
sulphuric  acid  continually  trickles  from  the  tank  at  the  top  of 
the  tower  L.  This  80  per  cent  acid  is  obtained  from  the  reser- 
voir at  the  bottom  of  the  Glover  tower,  from  which  place  it  is 
forced  through  a  lead  pipe  P  to  the  top  of  the  Gay-Lussac  tower. 
In  the  latter  the  80  per  cent  acid  dissolves  practically  all  the 
oxides  of  nitrogen,  and  the  residual  gases,  consisting  mainly  of 
nitrogen,  leave  the  top  of  the  tower  and  pass  into  a  large  chimney 
which  keeps  up  a  sufficient  draught.  The  acid  drawn  from  the 
bottom  of  the  Gay-Lussac  tower  is  thus  strongly  charged  with 
oxides  of  nitrogen.  It  is  the  purpose  of  this  tower  to  preserve 
these  oxides.  This  acid,  together  with  some  of  the  chamber 
acid,  is  used  again  in  the  Glover  tower  as  already  explained. 

The  acid  produced  in  the  chambers  is  known  as  "  chamber 
acid."  It.  is  about  60  to  70  per  cent  strong,  i.e.  of  specific  gravity 
of  about  1.5  to  1.6.  The  acid  may  be  further  concentrated  by 
evaporation  in  leaden  pans  to  78  per  cent.  Stronger  acid  than 
this  attacks  lead  too  much,  and  so  the  78  per  cent  acid  must  be 
further  concentrated  by  evaporation  either  in  cast-iron,  glass,  or 
platinum  vessels.  The  chamber  acid  is  commonly  used  directly 
in  the  manufacture  of  so-called  "superphosphate"  fertilizers, 


SULPHUR,    SELENIUM,   AND  TELLURIUM  199 

and  the  acid  from  the  bottom  of  the  Glover  tower  is  employed 
in  the  Le  Blanc  soda  process. 

The  concentrated  sulphuric  acid  on  the  market  has  a  specific 
gravity  of  from  1.83  to  1.84,  and  hence  contains  from  93  to  98 
per  cent  of  H2SO4.  In  making  concentrated  sulphuric  acid,  the 
contact  process  already  described  obviously  has  distinct  advan- 
tages, and  it  is  fast  taking  the  place  of  the  lead  chamber  method. 
The  latter  will,  however,  very  likely  continue  to  serve  to  pre- 
pare the  more  dilute  acid,  for  which  purpose  it  is  well  adapted. 

The  amount  of  sulphuric  acid  produced  in  the  world  annually 
is  over  four  million  tons.  The  material  is  used  in  making 
soda,  aniline  dyes,  fertilizers,  and  explosives  like  gun  cotton, 
nitro-powder,  and  dynamite.  Again,  it  is  used  in  storage 
batteries,  in  converting  starch  to  sugar  in  the  glucose  industries, 
in  refining  petroleum,  in  making  alum,  copper  sulphate,  and 
many  other  sulphates  that  are  used  in  medicine  and  in  the  arts. 

Properties  of  Sulphuric  Acid.  —  Sulphuric  acid  is  a  colorless, 
odorless,  heavy,  oily  liquid  of  specific  gravity  1.8384  at  15°. 
It  has  a  very  great  affinity  for  water,  with  which  it  unites  with 
great  evolution  of  heat.  For  this  reason  the  acid,  when  it  is  to 
be  diluted  with  water,  must  always  be  poured  gradually  into  an 
excess  of  water.  It  is  dangerous  to  proceed  in  the  reverse 
manner,  that  is  to  pour  the  water  into  the  acid,  for  the  great 
amount  of  heat  suddenly  liberated  is  very  apt  to  lead  to  explo- 
sions throwing  the  acid  out  of  the  container.  On  account  of 
its  powerful  affinity  for  water,  sulphuric  acid  exercises  a  destruc- 
tive action  upon  all  plant  and  animal  tissues,  for  it  abstracts 
hydrogen  and  oxygen  from  them  in  proportions  to  form  water, 
thus  leaving  a  dark  brown  or  black,  charred  mass  behind.  So 
wood,  sugar,  cork,  muscular  tissues,  etc.,  are  charred  by  sul- 
phuric acid.  When  sulphuric  acid  is  mixed  with  water  a  very 
appreciable  contraction  occurs  ;  thus  500  cc.  sulphuric  acid 
mixed  with  500  cc.  water  yield  a  mixture  that  has  a  volume  of 
971  cc.  On  account  of  its  affinity  for  water,  concentrated  sul- 
phuric acid  is  very  often  used  as  a  drying  agent  in  various 
chemical  operations,  particularly  in  drying  certain  gases  that 
are  not  affected  by  the  acid. 

The  commercial  sulphuric  acid  commonly  contains  lead  sul- 
phate, arsenic,  and  oxides  of  nitrogen.  By  distilling  it  from 
retorts  of  platinum  it  may  be  purified.  When  pure  anhydrous 


200  OUTLINES   OF   CHEMISTRY 

sulphuric  acid  (that  is,  H2SO4,  also  called  the  monohydrate 
because  it  is  H2O-SO3)  is  heated,  it  begins  to  fume  at  about 
150°  because  of  the  escape  of  SO3.  At  338°  the  acid  boils  and 
the  distillate  contains  1.5  per  cent  water. 

This  98.5  per  cent  acid  thus  has  a  constant  boiling  point  and 
cannot  be  further  concentrated  by  fractional  distillation.  At 
85  mm.  pressure,  pure  H2SO4  boils  without  decomposition  at 
145°-146°.  The  monohydrate  H2SO4  melts  at  +10°.  The 
crystals  are  colorless  and  may  be  freed  from  adhering  sulphuric 
acid  by  means  of  a  properly  constructed  centrifugal  machine. 

Sulphuric  acid  is  a  very  strong  dibasic  acid.  It  is  capable  of 
forming  acid  sulphates,  like  NaHSO4,  and  normal  sulphates,  like 
Na2SO4.  As  it  is  also  non- volatile  except  at  comparatively  high 
temperatures,  it  is  very  often  used  in  liberating  other  acids 
from  their  salts.  Besides  acting  as  an  acid,  sulphuric  acid  may 
also  play  the  role  of  an  oxidizing  agent  toward  many  substances. 
So  by  means  of  hydrogen  it  may  be  reduced  to  sulphurous  acid. 
When  the  metals  act  on  sulphuric  acid,  the  hydrogen  liberated 
reduces  the  acid  when  the  latter  is  used  in  concentrated  form, 
sulphates  and  sulphurous  acid  being  formed  simultaneously. 
The  sulphurous  acid  formed  may,  of  course,  be  reduced  still 
further.  Gold  and  platinum  do  not  act  on  sulphuric  acid. 
The  other  metals  react  with  it  under  certain  conditions,  forming 
sulphates.  Dilute  sulphuric  acid  acts  readily  on  some  metals, 
like  zinc  and  magnesium,  at  room  temperatures  liberating  the 
hydrogen,  as  was  pointed  out  when  the  latter  element  was 
studied.  Upon  other  metals,  like  copper  and  lead,  for  instance, 
sulphuric  acid  acts  but  slightly.  Even  hot,  fairly  concentrated 
sulphuric  acid,  as  we  have  seen,  does  not  attack  lead  much. 
This  is  due  in  part  to  the  fact  that  the  lead  sulphate  formed  is 
difficultly  soluble  in  sulphuric  acid  and  so  forms  a  protective 
coating  on  the  lead.  On  the  other  hand,  copper  acts  on  hot 
concentrated  sulphuric  acid,  forming  copper  sulphate  and  sul- 
phur dioxide.  By  means  of  hydrobromic  or  hydriodic  acid, 
sulphuric  acid  is  readily  reduced  to  sulphurous  acid  and  to 
hydrogen  sulphide.  The  sulphates  are  all  soluble  in  water 
except  the  sulphate  of  barium.  The  sulphates  of  lead,  stron- 
tium, and  calcium  are  sparingly  soluble  in  water.  As  a  rule, 
sulphates  are  not  as  soluble  as  chlorides  and  nitrates.  Sul- 
phates of  the  alkalies  are  quite  stable  at  high  temperatures. 


SULPHUR,   SELENIUM,   AND   TELLURIUM  201 

Sulphates  of  the  heavy  metals  decompose  at  high  temperatures, 
yielding  oxides  of  the  metals  and  sulphur  trioxide. 

Hydrates  of  Sulphuric  Acid.  —  Pure  H2SO4  is  commonly  called 
the  monohydrate,  as  stated  above.  When  one  molecule  of 
water  is  added  to  it,  it  forms  crystals  of  the  composition 
H2SO4-H2O  or  H4SO5,  which  melt  at  8°.  These  are  called 
the  dihydrate.  By  a  further  addition  of  a  molecule  of  water 
a  trihydrate  H2SO4-2H2O  or  H6SO6,  also  called  orthosulphuric 
or  normal  sulphuric  acid,  is  formed.  It  is  evident  that  it  may 
be  regarded  as  S(OH)6,  in  which  sulphur  is  combined  with  six 
hydroxyl  groups.  The  trihydrate  does  not  form  crystals, 
except  at  ver}^  low  temperatures.  Its  existence  is  largely 
based  upon  the  fact  that  it  represents  the  composition  of  the 
compound  formed  when  sulphuric  acid  arid  water  react  with 
maximum  contraction  of  volume.  There  are  no  salts  of  either 
H4SO5  or  H6SO6  known.  In  all  its  salts  sulphuric  acid  is 
distinctly  dibasic. 

Pyrosulphuric  Acid.  —  When  sulphur  trioxide  is  dissolved  in 
pure  sulphuric  acid,  pyrosulphuric  acid  or  disulphuric  acid 
H2S2OT  is  formed.  It  consists  of  crystals  that  melt  at  36°, 
and  is  sometimes  called  solid  sulphuric  acid.  This  acid  fumes 
strongly  in  the  air.  The  fuming  sulphuric  acid  of  commerce 
consists  of  sulphuric  acid  containing  varying  amounts  of  sul- 
phur trioxide  in  solution.  An  acid  containing  10  to  20  per 
cent  of  additional  SO3  in  solution  used  to  be  called  Nordhausen 
sulphuric  acid.  It  was  prepared  by  Basil  Valentine  at  Erfurt 
in  1450  by  heating  partially  dehydrated  sulphate  of  iron. 
From  pyrosulphuric  acid,  sulphur  trioxide  may  readily  be  pre- 
pared by  heating.  The  so-called  "oleum"  of  commerce  con- 
sists of  about  80  per  cent  SO3  and  20  per  cent  H2SO4.  It  is 
used  industrially.  The  salts  of  pyrosulphuric  acid  are  called 
the  pyrosulphates.  They  are  readily  prepared  by  heating  acid 
sulphates,  thus :  — 

KHSO4^l  K2S2O7  +  H2O. 

The  water  escapes  as  vapor.  On  moistening  the  pyrosulphate 
with  water,  the  acid  sulphate  is  again  obtained,  so  that  the 
above  reaction  is  reversible. 

Thiosulphates. — When  a  solution  of  a  sulphite  is  boiled  with 
sulphur,  a  thiosulphate  results  :  — 


202  OUTLINES   OF   CHEMISTRY 

We  may  look  upon  this  salt  as  sodium  sulphate  in  which  one 
oxygen  atom  is  replaced  by  a  sulphur  atom,  whence  the  name 
thiosulphate.  Sodium  thiosulphate  is  used  in  photography,  and 
in  commerce  it  is  frequently  called  hyposulphite  of  soda  or 
"  hypo."  These  names  are  not  in  accord  with  chemical  usage, 
since  the  salt  is  not  a  salt  of  an  acid  containing  less  oxygen 
than  sulphurous  acid  H2SO3.  By  treating  a  thiosulphate  with 
hydrochloric  acid,  the  chloride  of  the  metal,  sulphur,  sulphur 
dioxide,  and  water  are  formed,  thus  :  — 

.  Na2S2O3  +  2  HC1  =  2  NaCl  +  S  +  SO2  +  H2O. 

Thiosulphuric  acid  H2S2O3  is  not  known  in  the  free  state. 
Its  salts  are  very  common,  but  attempts  to  isolate  the  acid  fail 
because  it  decomposes  into  the  products  indicated  by  the  above 
equation. 

Persulphates.  —  By  electrolyzing  a  concentrated  solution  of 
acid  potassium  sulphate,  potassium  persulphate  KSO4  is  readily 
obtained.  Sodium  persulphate  may  be  similarly  prepared.  It 
is  used  in  photography.  Persulphuric  acid  HSO4  is  unstable. 
It  may  be  prepared  by  dissolving  its  anhydride,  S2O7,  sulphur 
peroxide,  in  water,  thus :  — 

S207  +  H20  =  2  HSO4. 

Sulphur  peroxide  or  heptoxide  was  formed  by  Berthelot  by  the 
action  of  the  silent  electric  discharge  on  a  mixture  of  sulphur 
dioxide  and  oxygen.  It  is  unstable,  and  but  little  is  known 
about  it.  Persulphuric  acid  is  formed  to  a  slight  extent  in  the 
lead  storage  cells,  in  which  sulphuric  acid  of  specific  gravity  1.2 
is  commonly  used. 

Polythionic  Acids.  —  Polythionic  acids  contain  more  than 
one  sulphur  atom.  Of  these  thiosulphuric  acid  H2S2O3  is  the 
simplest  example.  The  following  acids  are  known  :  — 

Thiosulphuric  Acid  H2S2O3,  forms  thiosulphates,  like  Na2S2O3. 

Dithionic  Acid  H2S2O6,  forms  dithionates,  like      Na2S2O6. 

Trithionic  Acid  H2S3OQ,  forms  trithionates,  like     Na2S3O6. 

Tetrathionic  Acid  H2S4O6,  forms  tetrathionates,  like  Na2S4O6. 

Pentathionic  Acid  H2S5O6,  forms  pentathionates,  like  Na2S5O6. 

With  the  exception  of  thiosulphuric  acid  (which  is  known 
only  in  form  of  salts),  the  free  acids  are  known  only  in  aqueous 


SULPHUR,    SELENIUM,   AND  TELLURIUM  203 

solutions  ;  and  even  in  these  they  readily  decompose.  The 
salts,  however,  are  as  a  rule  quite  stable. 

Thionyl  Chloride. — Thionyl  chloride  SOC12  is  formed  when 
phosphorus  pentachloride  acts  on  sulphur  dioxide,  or  on  a 
sulphite,  thus :  — 

PC15  +  SO2  =  POC13  +  SOC12. 
2  PC15  +  K2SO3  =  2  POC13  +  2  KC1  +  SOC12. 

It  is  a  colorless  liquid  of  very  pungent  odor.  It  fumes  in  the 
air  and  is  readily  decomposed  by  water,  thus :  — 

SOC12  +  H20  =  S02  +  2  HC1. 

Thionyl  chloride  boils  at  78°.  Its  specific  gravity  at  0°  is 
1.676.  It  may  be  regarded  as  SO2  with  one  oxygen  atom 
replaced  by  two  chlorine  atoms. 

Sulphuryl  Chloride.  —  This  compound  is  made  by  the  action 
of  equal  volumes  of  chlorine  and  sulphur  dioxide  on  each  other 
in  sunlight,  or  in  presence  of  a  little  camphor,  thus :  — 

S02+C12  =  S02C12. 

It  may  be  regarded  as  SO3  with  one  oxygen  atom  replaced  by 
two  chlorine  atoms.  It  is  a  colorless  liquid  of  very  pungent 
odor.  It  boils  at  70°,  and  has  a  specific  gravity  of  1.66  at  20°. 
In  contact  with  the  air  it  fumes  strongly.  By  addition  of  one 
gram-molecule  of  water  to  one  gram-molecule  of  sulphuryl 
chloride,  chlorsulphonic  acid  is  formed,  thus :  — 

S02C12  +  H20  =  S02 .  Cl  -  OH  +  HCL 

Chlorsulphonic  acid  SO2-C1- OH  may  be  regarded  as  sulphuric 
acid  SO2(OH)2  with  one  OH  group  replaced  by  chlorine. 
With  more  water,  chlorsulphonic  acid  decomposes,  thus  :  — 

S02.C1.0H  +  H20  =  S02(OH)2  +  HCL 

Selenium.  —  This  element  belongs  to  the  rarer  elements,  for 
though  it  is  fairly  widely  distributed  in  nature,  it  generally 
occurs  in  extremely  small  quantities.  It  has  been  found  in  the 
free  state  in  Mexico ;  but  it  occurs  mainly  in  combination  with 
metals  like  lead,  copper,  iron,  silver,  and  thallium.  Not  infre- 
quently it  is  present  in  small  amount  in  pyrites,  and  so  in 
roasting  the  latter  the  selenium  is  oxidized  to  selenium  dioxide 
which  is  carried  into  the  dust  flues  of  sulphuric  acid  factories, 


204  OUTLINES   OF   CHEMISTRY 

In  these  flues  there  is  also  deposited  some  free  selenium,  for  the 
latter  forms  when  hot  sulphur  acts  on  selenium  dioxide.  This 
gets  into  the  lead  chambers,  where  it  is  reduced  to  selenium 
by  the  action  of  sulphur  dioxide,  and  so  accumulates  in  the 
slime  at  the  bottom  of  the  chambers.  In  1817  Berzelius  dis- 
covered selenium  in  the  slime  of  the  lead  chambers  at  Gripsholm. 
He  named  the  element  selenium,  from  the  Greek  word  mean- 
ing moon,  because  of  its  similarity  to  tellurium,  which  is  named 
from  tellus,  the  earth. 

There  are  three  varieties  of  selenium :  (1)  a  red  amorphous 
precipitate  which  dissolves  in  carbon  disulphide  arid  separates 
from  the  latter  solution  in  form  of  (2)  red  monoclinic  crystals 
fusing  at  170-180°,  which  are  also  soluble  in  carbon  disul- 
phide ;  and  (3)  a  bluish  gray,  metallic  form  which  crystallizes 
in  the  hexagonal  system  and  is  insoluble  in  carbon  disulphide. 
This  metallic  form  conducts  electricity  slightly,  which  property 
may  be  increased  tenfold  by  exposure  to  light.  The  conduc- 
tivity depends  on  the  intensity  of  the  light. 

The  metallic  form  has  a  specific  gravity  of  4.8,  melts  at  217°, 
and  boils  at  680°.  The  atomic  weight  of  selenium  is  79.2,  and 
at  high  temperatures  the  molecular  weight  corresponds  to  the 
formula  Se2. 

Compounds  of  Selenium.  — These  are  similar  to  the  compounds 
of  sulphur.  So  hydrogen  selenide  may  be  formed  by  treating 
ferrous  selenide  with  hydrochloric  acid  :  — 

FeSe  +  HC1  =  FeCl2  +  H2Se. 

The  compound  H2Se  is  a  gas  that  has  the  smell  of  horseradish 
and  is  more  poisonous  than  hydrogen  sulphide.  The  aqueous 
solution  deposits  selenium  on  exposure  to  the  air  or  to  oxygen. 
With  the  exception  of  the  selenides  of  the  alkalies,  the  com- 
pounds of  the  metals  with  selenium  are  difficultly  soluble  in 
water. 

With  chlorine,  selenium  forms  selenium  monochloride  Se2Cl2 
and  selenium  tetrachloride  SeCl4.  The  former  is  a  dark, 
brownish  yellow  oil  and  the  latter  a  light  yellow  crystalline 
solid. 

Selenium  dioxide  SeO2  is  a  solid  formed  by  burning  selenium 
in  the  air.  It  is  the  only  oxide  of  selenium  known.  It  forms 
long  white  prismatic  crystals  that  sublime  at  about  300°. 


SULPHUR,   SELENIUM,   AND   TELLURIUM  205 

When  sulphur  and  selenium  dioxide  are  heated  together,  sul- 
phur dioxide  and  selenium  are  formed  :  — 
S  +  SeO2  =  SO2  +  Se. 

By  oxidizing  selenium  with  nitric  acid,  selenious  acid  H2SeO3 
is  produced.  By  means  of  sulphur  dioxide,  selenious  acid  is 
reduced  to  selenium  :  — 

H2Se03  +  2  S02  4-  H2O  =  2  H2SO4  +  Se. 

In  this  way  the  element  is  formed  in  the  lead  chambers  of  the 
sulphuric  acid  factories. 

When  SeO2  and  SeCl4  react  with  each  other,  they  form 
SeOCl2,  selenyl  chloride  :  — 

SeO2  +  SeCl4  =  2  SeOCl2. 

The  compound  melts  at  10°  and  boils  at  179°. 

Selenic  acid  H2SeO4  is  formed  by  oxidation  of  selenious  acid 
by  means  of  chlorine  :  — 

H2SeO3  +  H2O  +  C12  5*  2  HC1  +  H2SeO4. 

The  action  is  reversible,  for  selenic  acid  is  able  to  liberate 
chlorine  from  hydrochloric  acid.  Selenic  acid  is  thus  a  more 
powerful  oxidizing  agent  than  sulphuric  acid.  The  latter 
oxidizes  hydrobromic  acid,  but  not  hydrochloric  acid.  Pure 
selenic  acid  is  a  solid  melting  at  62°.  The  95  per  cent  solution 
is  a  thick,  oily  liquid  not  unlike  sulphuric  acid  in  appearance. 

When  hydrogen  sulphide  is  passed  into  a  solution  of  seleni- 
ous acid,  selenium  sulphide  SeS  is  precipitated.  It  is  yellow 
in  color  and  does  not  dissolve  in  ammonium  sulphide. 

Tellurium. — Tellurium  is  one  of  the  rare  elements.  It  has 
been  found  in  the  free  state,  and  also  in  the  form  of  tellurides 
in  combination  with  gold,  silver,  lead,  and  bismuth.  It  occurs 
in  Colorado,  California,  Hungary,  Brazil,  and  the  Liparian 
Islands.  It  is  a  brittle,  crystalline,  silvery  white  substance 
having  metallic  luster.  In  precipitated  amorphous  form  it  is 
a  black  powder.  In  metallic  form  it  conducts  heat  and  elec- 
tricity like  other  metals.  It  has  a  specific  gravity  of  6.26  and 
melts  at  455°.  Its  atomic  weight  is  127.5;  and  at  1400°,  its 
boiling  point,  the  vapor  density  corresponds  to  the  formula 
Te2.  Tellurium  was  discovered  in  1782  by  M  tiller  von  Reich- 
enstein,  whose  work  was  confirmed  by  Klaproth  in  1798.  The 
latter  called  the  element  tellurium,  from  tellus,  earth. 


206  OUTLINES   OF  CHEMISTRY 

Compounds  of  Tellurium.  —  By  the  action  of  hydrochloric 
acid  upon  zinc  telluride  ZnTe,  hydrogen  telluride  H2Te  is 
formed :  — 

ZnTe  +  2  HC1  =  ZnCla  +  H2Te. 

The  product  is  generally  contaminated  with  some  hydrogen, 
which  is  liberated  simultaneously.  Hydrogen  telluride  is"  a 
colorless,  poisonous  gas  of  disagreeable  odor.  It  is  combusti- 
ble and  fairly  soluble  in  water.  Its  aqueous  solutions  when 
in  contact  with  oxygen  or  air  gradually  deposit  tellurium. 
When  conducted  into  solutions  of  metallic  salts,  tellurides  of 
the  metals  are  in  general  precipitated.  Such  tellurides  may 
also  be  prepared  by  heating  metals  with  tellurium. 

With  chlorine,  tellurium  forms  tellurium  dichloride  TeCl2 
and  tellurium  tetrachloride  TeCl4.  These  are  formed  when 
chlorine  is  passed  over  hot  tellurium.  If  the  chlorine  is  in  large 
excess,  the  tetrachloride  is  formed  ;  if  less  chlorine  is  used,  the 
dichloride  forms  together  with  some  tetrachloride.  The  di- 
chloride is  a  black  crystalline  substance  melting  at  175°  and 
boiling  at  324°.  The  tetrachloride  forms  white,  shining  crys- 
tals that  melt  at  224°  and  boil  at  380°.  Both  chlorides  are 
decomposed  by  water.  It  is  to  be  noted  that  the  dichloride 
TeCl2  is  not  analogous  to  the  lower  chloride  of  sulphur,  which 
is  S2C12.  Tellurium  dibromide  TeBr2  and  tetrabromide  TeBr4 
have  also  been  prepared.  Tellurium  diiodide  TeI2  and  tellurium 
tetraiodide  TeI4  are  also  known. 

When  sulphur  trioxide  acts  on  tellurium,  tellurium  sulphur 
trioxide  TeSO3,  a  red  amorphous  solid,  forms,  which  on  heating 
is  decomposed  into  sulphur  dioxide  and  tellurium  monoxide  TeO. 
The  latter  is  a  black,  amorphous  substance,  which  on  heating 
yields  tellurium  dioxide  TeO2  and  tellurium. 

When  heated  in  the  air,  tellurium  is  oxidized  to  tellurium 
dioxide  TeO2.  This  is  a  white  crystalline  powder  which  is 
volatile  at  red  heat  (i.e.  at  higher  temperatures  than  tellurium 
itself)  and  difficultly  soluble  in  water.  By  means  of  nitric  acid, 
tellurium  may  be  oxidized  to  tellurous  acid  H2TeO3.  This  is  a 
feeble  acid  that  forms  a  white  powder  which  is  slightly  soluble 
in  water.  On  heating,  it  decomposes  into  water  and  tellurium 
dioxide.  With  strong  bases  it  forms  both  acid  and  normal 
tellurites,  like  KHTeO3  and  K2TeO3.  However,  towards 
strong  acids  it  behaves  like  a  base.  The  salts  thus  formed 


SULPHUK,    SELENIUM,   AND   TELLURIUM  207 

may  be  considered  as  derivatives  of  Te(OH)4,  that  is, 
H2TeO3-H2O.  So,  for  instance,  tellurium  sulphate  Te(SO4)2 
has  been  prepared.  Moreover,  the  salt  TeCl4  may  be  retained  in 
aqueous  solutions  in  presence  of  an  excess  of  hydrochloric  acid. 
Being  both  a  weak  base  and  also  a  weak  acid,  the  salts  that  tel- 
lurous  acid  forms  with  either  bases  or  acids  are  not  very  stable. 
This  is  generally  the  case  with  substances  that  do  not  have  pro- 
nounced chemical  characteristics. 

On  fusing  together  barium  nitrate  and  tellurium  dioxide, 
barium  tellurate  may  be  formed:  — 

Ba(NO3)2  +  TeO2  =  BaTeO4  +  2  NO2. 

By  decomposing  barium  tellurate  with  the  calculated  quantity 
of  sulphuric  acid,  barium  sulphate,  which  is  insoluble  in  water, 
and  telluric  acid  H2TeO4,  which  remains  in  solution,  result:  — 

BaTeO4  +  H2SO4  =  BaSO4  +  H2TeO4. 

The  latter  may  also  be  prepared  by  first  making  potassium  tel- 
lurate K2TeO4,  by  fusing  either  tellurium  or  tellurium  dioxide 
with  potassium  carbonate  and  potassium  nitrate,  or  by  passing 
chlorine  into  an  alkaline  solution  of  potassium  tellurite.  The 
potassium  tellurate  is  then  converted  into  the  barium  salt  by 
means  of  barium  chloride,  thus:  — 

K2Te04  +  BaCl2  =  2  KC1  +  BaTeO4  ; 

and  the  barium  tellurate  is  then  decomposed  by  dilute  sul- 
phuric acid  as  before.  From  the  aqueous  solution,  telluric  acid 
separates  in  form  of  monoclinic  crystals  of  the  composition 
H2TeO4-2  H2O  or  Te(OH)6.  On  heating  these,  H2TeO4  forms, 
which  loses  water  at  160°,  yielding  tellurium  trioxide  TeO3,  an 
orange-yellow,  crystalline  substance  that  unites  with  water  ex- 
tremely slowly  and  decomposes  into  tellurium  dioxide  and 
oxygen  on  ignition.  While  telluric  acid  forms  tellurates  with 
the  alkalies  and  other  metals,  its  resemblance  to  sulphuric  acid 
and  selenic  acid  is  extremely  slight.  Like  tellurous  acid,  telluric 
acid  may  act  as  a  base  toward  strong  acids. 

General  Considerations.  —  Oxygen  is  commonly  considered 
as  forming  with  sulphur,  selenium,  and  tellurium  a  natural 
family  group  of  elements.  We  have  already  seen  that  fluorine, 
chlorine,  bromine,  and  iodine  form  such  a  group  in  which 


208 


OUTLINES   OF   CHEMISTRY 


fluorine  is  rather  less  closely  related  to  chlorine,  bromine,  and 
iodine,  than  these  three  are  to  one  another.  Now,  the  relation 
of  oxygen  is  similarly  less  close  to  sulphur,  selenium,  and  tel- 
lurium. From  oxygen  to  tellurium  we  have  a  gradation  of 
physical  properties,  as  the  following  table  shows  :  — 


NAME  / 

COLOR 

ATOMIC 
WEIGHT 

SPECIFIC  GRAVITY 

MELTING 
POINT 

BOILING 
POINT 

Oxygen 

blue 

16.0 

1.124  (at  -  181°) 

-181.4 

Sulphur 

yellow 

32.07 

1.96  to  2.0 

114.5 

450 

Selenium 

red  or 

metallic 

79.2 

4.8 

217 

680 

Tellurium 

black  or 

metallic 

127.5 

6.3 

455 

1400 

All  of  these  elements  exhibit  allotropism. 
Toward  hydrogen  these  elements  are  bivalent,  forming  com- 
pounds of  the  type  H2X,  thus :  — 

H20,  H2S,  H2Se,  H2Te. 

The   stability   of    these   compounds   decreases   as   the   atomic 
weight  of  the  elements  in  question  increases. 

Sulphur,  selenium,  and  tellurium  form  compounds  with 
oxygen,  whose  composition  is  represented  by  the  types  XO2  and 
XO3.  In  the  former,  that  is  SO2,  SeO2,  and  TeO2,  the  elements 
are  tetravalent ;  whereas  in  the  latter,  namely,  SO3,  SeO3,  and 
TeO3,  the  elements  in  question  are  hexavalent,  which  is  the  high- 
est valence  they  are  capable  of  exhibiting.  Again,  in  the  acids 
of  the  type  H2XO3,  namely  H2SO3,  H2SeO3  and  H2TeO3,  and 
in  those  of  the  type  H2XO4,  namely  H2SO4,  H2SeO4,  and 
H2TeO4,  we  plainly  have  striking  analogies.  In  the  com- 
pounds H2XO3,  the  elements  are  quadrivalent,  thus  :  — 

X)  -  H  /0-K  /0-K 

84=0          ,         Se^O          ,        Te==0 
\0  -  H  \0  -  H  X)  -  H 

In  the  compounds  H2XO4,  the  elements  are  hexavalent,  thus :  — 


0-H 


O-H 


O-H 


O-H. 


)-H 


SULPHUR,    SELENIUM,    AND   TELLURIUM  209 

Toward  the  halogens,  sulphur,  selenium,  and  tellurium  are 
bivalent  and  quadrivalent,  while  in  some  oxy -halogen  deriva- 
tives they  are  hexavalent,  thus  :  — 


Cl 
Cl 

Cl  /Cl 


s^c!  Se<ci 


VX3L- 

\C1 


/Cl 


Cl  \C1 

/f 


s\r°'       * 

\C1  XC1  \^ 

The  halogen  compounds  of  selenium  are  more  stable  than 
those  of  sulphur,  and  those  of  tellurium  are  more  stable  than 
the  selenium  halides.  One  may  regard  the  very  unstable  com- 
pound C12O  as  analogous  to  C12S,  Cl2Se,  and  Cl2Te.  Thus,  it 
is  apparent  that  as  the  atomic  weight  of  the  elements  of  the 
oxygen  group  increases,  their  affinity  for  halogen  increases. 
This  is  also  evident  from  the  fact  that  in  the  halogen  com- 
pounds, sulphur  is  readily  replaced  by  selenium,  and  the  latter 
is  in  turn  replaced  by  tellurium. 

Sulphur,  selenium,  and  tellurium,  like  crhlorine,  bromine,  and 
iodine,  form  a  group  of  three,  in  which  the  atomic  weight  of 
the  middle  element  is  very  nearly  equal  to  one-half  the  sum  of 
the  atomic  weights  of  the  other  two,  thus  :  — 

S  Te 

J(32.07 +  127.5)  =79.78;  whereas,  Se  =  79.2. 

In  spite  of  the  relationships  noted,  it  should  be  borne  in 
mind,  however,  that  tellurium  after  all  shows  some  decided 
points  of  departure  in  its  chemical  behavior  from  sulphur  and 
selenium,  so  that  the  closeness  of  relationship  between  the 
latter  elements  and  tellurium  has  repeatedly  been  called  into 
question. 


CHAPTER  XIV 

CARBON   AND   SOME   OF   ITS   TYPICAL   COMPOUNDS 

Occurrence  and  Allotropic  Forms  of  Carbon. — Carbon  occurs 
in  the  free  state  in  nature  as  diamond  and  graphite,  which  are 
crystalline  in  character.  It  is  also  known  in  amorphous  form 
as  charcoal,  coke,  soot,  lampblack,  bone  black,  etc.,  resulting 
from  the  charring  of  animal  and  vegetable  matter,  and  various 
compounds  of  carbon.  Diamond,  graphite,  and  amorphous 
carbon  are  the  allotropic  forms  of  the  element. 

Carbon  is  a  most  important  constituent  of  all  plants  and 
animals.  Large  quantities  of  carbon  are  found  in  form  of  coal, 
which  represents  the  remains  of  vegetation  of  past  geological 
ages.  Coal  is  consequently  not  pure  carbon.  Indeed,  the 
amount  of  free  carbon  in  different  kinds  of  coal  varies  con- 
siderably. In  natural  gas  and  petroleum,  carbon  occurs  in 
combination  with  hydrogen.  In  form  of  carbon  dioxide,  carbon 
is  found  in  the  air  and  in  many  natural  waters.  As  carbonates, 
especially  calcium  carbonate  and  magnesium  carbonate,  carbon 
is  found  in  huge  masses  widely  distributed  in  various  parts  of 
the  earth.  Calcium  carbonate  is  commonly  found  in  form  of 
chalk,  limestone,  and  marble ;  whereas  calcium  magnesium 
carbonate,  or  dolomite,  also  called  magnesian  limestone,  occurs 

very    frequently    and     covers 
extensive  areas  of  the  earth's 
crust,  often  forming  mountains. 
Diamond.  —  The  diamond  is 
a  crystalline  form  of  carbon, 
and  belongs  to  the  regular  or 
FlG-  8L  FlG-  82-  isometric  system,  Figs.  81  and 

82.  It  is  colorless  when  pure,  but  frequently  it  is  dark-colored 
or  black,  in  which  form  it  is  known  as  carbonado,  or  bort.  Some- 
times diamonds  are  colored  blue,  green,  yellow,  or  red  by  small 
amounts  of  foreign  substances.  The  diamond  is  very  hard. 
It  will  scratch  all  other  minerals.  For  this  reason  diamond 

210 


.  CARBON   AND   SOME-  OF  ITS  TYPICAL   COMPOUNDS        211 

dust,  usually  in  form  of  bort,  is  used  by  lapidaries  in  cutting 
and  polishing.  Drills  of  carbonado,  which  sometimes  occurs 
in  pieces  as  large  as  one's  fist,  are  used  in  boring  rocks,  and 
glaziers  make  use  of  the  diamond  in  cutting  glass.  Diamonds 
are  found  in  Brazil,  South  Africa,  India,  and  Australia,  and 
sometimes  also  in  the  United  States.  They  are. usually  cor- 
roded, and  so  their  brilliant  luster  does  not  appear  till  the  outer 
layer  is  removed  by  the  lapidary,  who,  by  grinding  suitable 
artificial  faces  on  the  diamonds,  brings  out  their  highly  prized 
brilliancy.  The  diamond  has  a  specific  gravity  of  3.5  to  3.6 
and  a  refractive  index  of  2.416  to  2.43.  That  it  consists  of 
carbon  only,  appears  from  the  fact  that  on  combustion  in  oxy- 
gen the  sole  product  formed  is  carbon  dioxide.  Diamonds  of 
microscopic  dimensions  have  been  prepared  artificially  by  Mois- 
san,  by  dissolving  carbon  in  molten  iron  and  then  chilling  the 
same  suddenly.  Under  the  great  pressure  thus  produced  in  the 
interior  of  the  iron,  graphite  and  some  very  small  diamonds 
crystallize  out.  The  diamond  is  not  attacked  by  acids.  It  is  a 
non-conductor  of  electricity,  and  becomes  electrically  charged 
when  rubbed  with  a  cloth.  On  heating  it  highly  out  of  contact 
with  the  air,  it  may  be  changed  to  graphite ;  whereas  when 
heated  highly  in  oxygen,  it  burns  with  great  brilliancy  to  car- 
bon dioxide. 

The  diamonds  found  are  usually  small ;  only  rarely  do  they 
weigh  as  much  as  20  grams.  The  largest  diamond  known  was 
found  near  Pretoria,  South  Africa.  It  is  called  the  Cullinan, 
and  weighed  about  600  grams  when  found. 

Graphite.  —  Graphite  crystallizes  in  the  monoclinic  system  in 
plates  that  simulate  hexagonal  forms.  It  is  widely  distributed 
in  nature  in  form  of  small  flakes  or  granules  in  various  granite 
rocks.  It  is  black  or  grayish  black,  having  metallic  luster.  It 
is  very  soft,  and  may  readily  be  crushed  to  a  fine  powder,  which 
is  frequently  employed  as  a  lubricant.  Graphite  is  also  used 
in  making  ulead  pencils."  At  one  time  it  was  thought  that 
graphite  contained  lead,  hence  the  name  plumbago,  by  which 
graphite  is  also  known.  Graphite  conducts  electricity,  and  in 
its  artificial  forms  it  is  used  as  electrodes  for  electric  arcs  and 
in  electro-chemical  work,  especially  in  making  chlorine  and  caus- 
tic soda  from  common  salt.  Graphite  powder  is  employed  in 
electrotyping,  in  making  stove  polish,  etc.  The  specific  gravity 


212  OUTLINES   OF  CHEMISTRY 

of  graphite  varies  from  1.8  to  2.5.  It  burns  with  great  difficulty 
even  in  oxygen,  and  the  natural  varieties  leave  from  3  to  20 
per  cent  ash,  which  usually  consists  of  silicates  of  various  bases. 
Graphite  is  very  refractory  and  not  readily  attacked  by  chemi- 
cal reagents,  for  this  reason  it  is  frequently  employed  together 
with  fire  clay  in  making  crucibles.  A  mixture  of  concentrated 
nitric  acid  and  potassium  chlorate  converts  graphite  into  gra- 
phitic acid,  which  consists  of  small  yellow  crystals  that  explode 
on  heating,  leaving  a  mass  of  finely  divided  carbon.  Graphitic 
acid  consists  of  56  per  cent  carbon,  2  per  cent  hydrogen,  and 
42  per  cent  oxygen.  On  treating  graphite  with  concentrated 
nitric  acid,  and  then  igniting  the  material  strongly,  various 
samples  show  a  different  behavior.  Thus,  graphite  found  in 
the  State  of  New  York  greatly  increases  in  volume  when  so 
treated,  whereas  Siberian  graphite  is  not  so  affected'  at  all. 
Ceylon  and  Siberia  furnish  most  of  the  natural  graphite.  Arti- 
ficial graphite  is  formed  when  carbon  dissolved  in  molten  iron 
crystallizes  out,  and  when  coke  is  heated  to  very  high  tempera- 
tures in  the  electric  furnace  out  of  contact  of  the  air  and  then 
allowed  to  cool  slowly.  By  the  latter  process,  known  as  the 
Acheson  process,  much  graphite  of  excellent  quality  is  made  at 
Niagara  Falls.  By  first  grinding  up  the  coke,  mixing  it  with 
a  binder,  like  coal  tar  or  black  strap  molasses,  and  molding  it 
into  desired  forms  and  baking  these  in  ovens,  carbons  for  bat- 
teries, electric  arc  lights,  and  other  purposes  are  obtained. 
These  carbons  may  readily  be  converted  into  graphite  by  heat- 
ing them  in  the  electric  furnace  as  above  stated.  Artificial 
graphite  is  now  much  employed  in  the  arts  in  place  of  natural 
graphite. 

Electric  furnaces  are  either  resistance  furnaces  or  arc  furnaces. 
In  the  former  the  heat  is  generated  by  passing  a  strong  current 
of  electricity  through  conducting  material  which  is  thus  heated 
to  high  temperatures.  In  the  arc  furnaces,  the  high  tempera- 
tures are  obtained  by  means  of  the  electric  arc ;  in  principle, 
the  arrangement  is  similar  to  that  of  the  arc  lamp.  The 
Acheson  graphite  furnace  is  a  form  of  resistance  furnace, 
Fig.  83.  The  molded  sticks  of  carbon  are  piled  between  two 
carbon  terminals  which  are  about  two  feet  square  and  thirty 
feet  apart.  The  whole  is  then  covered  with  a  thick  layer  of 
granular  carbon  and  carborundum  (which  see),  and  a  current 


CARBON  AND  SOME   OF  ITS  TYPICAL   COMPOUNDS        213 


of  3000  amperes  at  a  pressure  of  220  volts  is  turned  on  which 
is  gradually  changed  to  9000  amperes  at  80  volts  at  the  end  of 


FIG.  83. 

the  twenty-four  hours  during  which  the  furnace  is  run.  The 
whole  is  then  allowed  to  cool,  and  the  carbon  is  found  to  be 
converted  to  graphite.  The  heat  is  largely  generated  by  the 
current  as  it  passes  through  the  granular  carbon  lodged  in  the 
interstices  between  the  carbon  sticks.  Smaller  resistance 
furnaces  of  different  types  are  used  in  the  arts  and  also  for 
experimental  purposes. 

Figure  84  shows  a-  typical  arc  furnace  for  experimental  work. 
The  poles  are  of  carbon,  and  the  walls  of  the  furnace  are  usually 
built  of  quicklime  or  mag- 
nesia. In  the  case  of  larger 
furnaces,  this  material  is  then 
reenforced  on  the  outside  by 
bricks  made  of  crude  mag- 
nesia. We  shall  have  occa- 
sion to  refer  to  electric  fur- 
naces again  as  we  proceed, 
for  they  are  now  used  for  various  purposes  that  require  high 
temperatures. 

Amorphous  Carbon.  —  Amorphous  carbon  is  formed  when 
animal  or  vegetable  matter,  coal,  or  various  compounds  of  car- 
bon are  heated  while  access  of  air  is  excluded  entirely  or  in 
part.  Thus,  by  heating  wood  in  this  way  charcoal  is  produced. 


FIG.  84. 


214 


OUTLINES   OF   CHEMISTRY 


Similarly,  bone  black  is  made  by  heating  bones  in  closed  iron 
cylinders ;  coke  is  formed  by  heating  coal ;  and  blood  charcoal 
is  produced  by  heating  blood  till  it  chars.  On  burning  hydro- 
carbons, like  kerosene  or  turpentine,  lampblack,  or  soot,  is 
deposited  on  a  cold  object  held  in  the  smoky  flame.  Lamp- 
black is  nearly  pure  carbon.  All  of  these  forms  of  amorphous 
carbon  differ  according  to  the  source  from  which  they  are 
obtained. 

Charcoal  contains  about  85  per  cent  carbon.  It  is  very 
porous  and  when  freshly  heated  it  has  the  power  of  absorbing 
many  gases.  So  freshly  ignited  charcoal  absorbs  from  50  to 
100  times  its  own  volume  of  ammonia,  hydrogen  sulphide, 

bromine  vapor,  etc.,  which  are  again 
given  off  on  heating  the  charcoal. 
This  process  of  absorption  consists  of 
condensation  of  the  gases  on  the  sur- 
face of  the  charcoal  and  is  called 
adsorption.  Figure  85  illustrates  th^ 
adsorption  of  ammonia  (which  has  been 
collected  over  mercury)  by  means  of 
charcoal.  Charcoal  is  frequently  used 
as  a  deodorant  of  vaults,  cisterns,  etc., 
because  it  absorbs  the  noxious  gases. 
Bone  black  acts  similarly.  The  latter 
is  much  used  to  decolorize  various 
solutions.  So  in  the  refining  of  sugar, 
bone  black  serves  to  remove  the  brown 
coloring  matter.  Bone  black  contains 
from  70  to  80  per  cent  calcium  phosphate,  and  only  from  8  to  12 
per  cent  carbon,  together  with  some  calcium  carbonate  and 
calcium  sulphate.  By  repeated  treatment  with  acids  and  water, 
nearly  pure  carbon  may  be  prepared  from  bone  black.  A  very 
pure  amorphous  carbon  may  be  prepared  by  charring  sugar. 
All  animal  charcoal  contains  some  nitrogen,  which  is  very 
tenaciously  held.  Very  finely  divided,  dry  charcoal  has  a 
strong  affinity  for  oxygen ;  indeed,  it  may  catch  fire  on  being 
thrown  from  its  container  into  the  open  air,  —  pyrophoric  carbon. 
The  various  forms  of  amorphous  carbon  differ  in  their  physical 
nature,  and  in  the  chemical  character  and*  the  amounts  of  the 
impurities  they  contain. 


FIG.  85. 


CARBON  AND   SOME   OF  ITS  TYPICAL   COMPOUNDS        215 

Lampblack  is  much  used  as  a  pigment  in  paints,  India  ink, 
etc.,  whereas  coke  is  used  for  fuel  and  in  the  reduction  of  ores, 
particularly  iron  ore.  In  the  production  of  coke,  the  gaseous 
products,  tar,  etc.,  should  be  saved  and  not  allowed  to  burn 
and  go  to  waste  as  is  still  frequently  done.  Similarly  many 
valuable  volatile  products  are  lost  when  charcoal  is  made  by 
heating  wood  in  a  pile  covered  with  sod  and  earth  in  the  old- 
fashioned  way.  By  heating  wood  and  coal  in  retorts,  these 
gaseous  and  other  volatile  products  may  be  saved  and  used. 
The  process  of  thus  heating  and  decomposing  substances  in 
retorts  is  called  destructive  or  dry  distillation. 

Coal.  —  Coal  is  found  in  large  deposits  in  various  parts  of 
the  earth.  It  represents  the  plant  remains  of  various  geologi- 
cal periods  from  the  carboniferous  to  the  tertiary.  Coal  has 
been  formed  by  the  gradual  abstraction  of  carbon,  hydrogen, 
and  oxygen,  mainly  in  form  of  water,  marsh  gas,  etc.,  from  the 
vegetable  remains.  Much  marsh  gas  (which  see)  is  found 
associated  with  coal.  There  are  many  varieties  of  coal,  which 
are  commonly  roughly  divided  into  two  great  classes,  namely, 
soft  coal  and  hard  coal.  Hard  coal,  or  anthracite,  contains  much 
less  volatile  matter  than  soft  coal,  which  is  also  called  bitu- 
minous coal.  The  latter  burns  with  a  sooty  flame,  and  evolves 
more  hydrocarbon  gases  than  anthracite  when  heated.  It  is 
consequently  used  in  manufacturing  illuminating  gas  (which 
see).  The  amount  of  carbon  in  charcoal,  coke,  or  anthracite 
is  generally  in  the  neighborhood  of  95  per  cent,  whereas  soft 
coal  contains  about  80  per  cent  carbon. 

Chemical  Behavior  of  Carbon.  —  The  various  forms  of  carbon 
above  mentioned  are  practically  not  at  all  attacked  by  chemi- 
cal reagents  at  room  temperatures.  When  heated  in  the  air, 
carbon  burns ;  that  is,  it  unites  with  oxygen.  When  a  suffi- 
cient supply  of  oxygen  is  at  hand,  the  product  formed  is  car- 
bon dioxide,  a  gas  whose  composition  corresponds  to  the 
formula  CO2.  When  an  insufficient  amount  of  oxygen  is 
furnished,  some  carbon  monoxide  is  also  formed.  The  latter 
substance  is  also  a  gas ;  its  composition  is  expressed  by  the 
formula  CO. 

The  very  careful  work  of  Dumas  and  Stas  has  shown  that  3 
parts  of  carbon  by  weight  unite  with  8  parts  of  oxygen  by 
weight  to  fornv  carbon  dioxide  gas.  Under  standard  conditions 


216  OUTLINES   OF   CHEMISTRY 

the  weight  of  22.4  liters  of  carbon  dioxide  is  nearly  44  grams. 
The  molecular  weight  of  carbon  dioxide  is  consequently  44. 
However,  44  parts  of  carbon  dioxide  by  weight  contain  12 
parts  of  carbon  and  32  parts  of  oxygen.  Now,  since  we  have 
previously  chosen  16  as  the  atomic  weight  of  oxygen,  it  is  clear 
that  the  molecule  of  carbon  dioxide  contains  32-^16  or  2  atoms 
of  oxygen.  The  question  now  arises,  how  many  carbon  atoms 
there  are  in  a  molecule  of  carbon  dioxide  ?  To  answer  this 
question  really  involves  choosing  the  atomic  weight  of  carbon 
from  the  combining  weight.  A  study  of  all  of  the  gaseous 
compounds  of  carbon  that  are  known  has  revealed  the  fact  that 
in  no  case  do  these  contain  less  than  12  grams  of  carbon  in 
22.4  liters  under  standard  conditions  ;  that  is  to  say,  there  is 
no  compound  into  which  carbon  enters  whose  molecule  con- 
tains less  than  12  parts  of  carbon  by  weight.  For  this  reason, 
the  atomic  weight  of  .carbon  is  taken  as  12,  rather  than  some 
fraction  or  multiple  thereof.  Having  thus  determined  upon 
12  as  the  atomic  weight  of  carbon,  it  is  obvious  from  what  has 
been  said  that  the  molecule  of  carbon  dioxide  contains  one 
atom  of  carbon  and  two  atoms  of  oxygen,  and  its  formula  is 
therefore  CO2.  The  specific  heat  of  carbon  increases  with  the 
temperature,  becoming  nearly  constant  in  the  neighborhood  of 
1000°.  Between  900°  and  1000°  it  is  about  0.46,  which  fact 
also  yields  approximately  12  as  the  atomic  weight  of  carbon 
according  to  the  law  of  Dulong  and  Petit. 

When  carbon  is  burned  in  oxygen,  the  volume  of  the  car- 
bon dioxide  formed  is  the  same  as  that  of  the  original  oxygen, 
measured,  of  course,  under  the  same  conditions  of  temperature 
and  pressure.  Figure  77  represents  an  apparatus  for  demon- 
strating this  fact.  Carbon  on  a  platinum  spoon  is  burned 
in  the  flask  which  is  filled  with  oxygen.  After  the  whole  has 
again  cooled  to  room  temperature,  the  mercury  manometer, 
attached  as  shown,  indicates  that  there  has  been  no  change  of 
volume.  Therefore  we  have  :  — 


(1  volume)  (1  volume) 

which  is  quite  in  harmony  with  what  we  should  expect  on  the 
basis  of  Avogadro's  hypothesis.  It  will  be  recalled  that  when 
sulphur  is  similarly  burned  in  oxygen  the  volume  of  the  SOa 
formed  is  also  the  same  as  that  of  the  oxygen.  t 


CARBON  AND  SOME  OF  ITS  TYPICAL  COMPOUNDS        217 

Carbon  does  not  unite  directly  with  hydrogen  except  at  very 
high  temperatures  such  as  are  produced  by  the  electric  arc,  and 
even  then  the  union  takes  place  with  difficulty.  Very  many 
compounds  of  hydrogen  and  carbon  are  known,  however,  for 
they  may  be  prepared  by  indirect  methods.  Similarly  carbon 
does  not  combine  directly  with  the  halogens,  except  in  the  case 
of  fluorine.  However,  by  indirect  methods  compounds  of  car- 
bon with  the  halogens  may  be  formed  fairly  readily.  With 
sulphur  carbon  unites  directly  at  high  temperatures.  At  the 
temperature  of  the  electric  furnace,  carbides  of  calcium,  alumi- 
num, silicon,  and  boron  may  be  formed,  and  iron  generally 
contains  some  carbon  which  is  present  in  form  of  carbide  of 
iron. 

With  nitrogen  carbon  does  not  unite  directly,  though  by 
indirect  means  it  is  quite  possible  to  effect  the  union  of  these 
two  elements. 

In  general,  carbon  is  rather  inert  chemically,  though  its  com- 
pounds, when  once  formed,  frequently  have  a  very  considerable 
degree  of  stability.  We  shall  see  that  the  carbon  atom  has  a  great 
tendency  to  unite  with  other  carbon  atoms,  which  often  results  in 
building  up  of  large  and  complex  molecules.  At  high  tempera- 
tures carbon  is  generally  bivalent,  whereas  at  lower  tempera- 
tures it  is  quadrivalent.  Sometimes  carbon  is  regarded  as 
quadrivalent  in  all  of  its  compounds.  This  view  would  neces- 
sitate that  a  goodly  number  of  carbon  compounds  be  regarded 
as  unsaturated.  The  number  of  carbon  compounds  known  is 
very  great,  so  that  it  is  customary  to  treat  these  in  a  separate 
division  of  chemistry,  namely,  organic  chemistry.  The  term 
organic  chemistry,  as  used  at  present,  signifies  that  branch  of  the 
science  which  deals  with  the  compounds  of  carbon,  or  with  the 
hydrocarbons  and  their  derivatives  ;  for  all  the  carbon  compounds 
may  be  considered  as  derived  from  compounds  of  carbon  and 
hydrogen  by  substituting  other  elements  in  place  of  the  hydrogen. 
Since  carbon  is  an  essential  constituent  of  all  living  beings,  the 
study  of  the  carbon  compounds  is  closely  associated  with  the 
study  of  the  chemistry  of  the  products  that  are  formed  in  or- 
ganic beings.  Hence  the  name  organic  chemistry.  Indeed, 
till  1828  it  was  quite  generally  held  that  compounds  that  are 
produced  by  the  life  process  could  not  be  prepared  artificially 
in  the  laboratory,  but  Wohler's  synthesis  of  urea,  (which  see) 


218  OUTLINES   OF  CHEMISTRY 

showed  that  such  syntheses  are  quite  possible ;  and  now  many 
of  the  products  that  are  formed  by  the  metabolic  processes  in 
plants  and  animals  have  been  prepared  in  the  laboratory. 

Carbon  Dioxide.  —  As  already  mentioned,  there  are  two  oxides 
of  carbon  known,  namely,  carbon  monoxide  CO,  and  carbon 
dioxide  CO2.  Carbon  dioxide  is  the  final  oxidation  product 
of  carbon.  It  always  results  when  carbon  is  burned  in  a  suffi- 
cient amount  of  air  or  oxygen.  It  is  moreover  exhaled  as  a 
product  of  respiration  of  both  plants  and  animals.  It  will  be 
recalled  that  every  1000  volumes  of  air  contain  3  volumes  of  car- 
bon dioxide.  This  gas  is  also  contained  dissolved  in  all  natural 
waters.  Many  spring  waters,  like  those  at  Saratoga,  Colorado 
Springs,  Selters,  Vichy,  and  Narzan,  are  so  highty  charged  with 
carbon  dioxide,  under  pressure,  that  the  gas  escapes  with  effer- 
vescence when  the  pressure  is  released.  In  volcanic  regions, 
like  those  of  Italy,  South  America,  and  Java,  large  quantities 
of  carbon  dioxide  issue  from  fissures  in  the  earth's  crust.  In 
fermentation  and  in  the  decay  of  animal  and  vegetable  matter 
carbon  dioxide  is  always  formed.  It  is  thus  evident  that  the 
air  is  continually  receiving  carbon  dioxide  from  quite  a  variety 
of  sources. 

Carbon  dioxide  maybe  produced  by  the  oxidation  of  carbon :  — 

C  +  02=C02. 

In  the  process  of  fermentation  of  sugar  by  means  of  yeast, 
alcohol  and  carbon  dioxide  result.  Thus  the  fermentation  of 
glucose  may  be  expressed  by  the  following  equation :  — 

C6H1206=2C2H6OH  +  2C02. 

glucose  alcohol          carbon  dioxide 

However,  the  simplest  way  of  obtaining  carbon  dioxide  is  by 
the  action  of  an  acid  upon  a  carbonate,  like  calcium  carbonate 
CaCO3,  or  sodium  carbonate  :  — 

CaCO3  +  2  HC1  =  CaCl2  +  H2O  +  CO2. 
Na2C03  +  H2S04  =  Na2S04  +  H2O  +  CO2. 

It  will  be  recalled  that  when  an  acid  acts  on  a  sulphite,  H2SO3 
is  not  formed,  for  it  at  once  decomposes,  yielding  H2O  and  SO2. 
Similarly,  when  an  acid  acts  on  a  carbonate,  we  do  not  get 
H2CO3,  carbonic  acid,  but  H2O  and  CO2.  The  latter  is  clearly 
carbonic  acid  anhydride.  It  is  probable  that  the  compound 


CARBON  AND   SOME    OF   ITS  TYPICAL   COMPOUNDS        219 


H2SO3,  sulphurous  acid,  exists  in  aqueous  solutions  of  SO2; 
and  it  is  also  likely  that  aqueous  solutions  of  CO2  contain 
H2CO3,  carbonic  acid.  While  an  aqueous  solution  of  any  of 
the  ordinary  acids  will  decompose 
carbonates,  owing  to  the  weak- 
ness of  carbonic  acid  and  the 
volatility  of  carbon  dioxide,  cal- 
cium carbonate,  in  form  of  marble, 
and  hydrochloric  acid  are  com- 
monly employed,  because  these 
materials  are  cheap,  and  the  cal- 
cium chloride  is  readily  soluble 
in  water.  A  Kipp  apparatus 
(Fig.  86)  is  very  convenient  for 
evolving  carbon  dioxide  from 
marble  and  hydrochloric  acid. 

Just  as  sulphites  are  salts  of 
sulphurous  acid  H2SO3,  so  car- 
bonates may  be  regarded  as  salts 
of  the  dibasic  acid  H2CO3.  We 
have  then  normal  carbonates,  like 
Na2CO3,  and  acid  carbonates,  or 
bicarbonates,  like  NaHCO3.  Basic  carbonates  of  many  metals, 
like  zinc,  copper,  lead,  etc.,  are  also  known. 

When  carbon  dioxide  is  conducted  into  clear  limewater,  a 
white  precipitate  of  calcium  carbonate  is  formed,  thus:  — 

Ca(OH)2  +  G02  =  CaC03  +  H2O. 

On  continuing  to  conduct  in  more  carbon  dioxide  the  calcium 
carbonate  again  dissolves,  the  solution  becoming  clear.  In 
this  process  calcium  bicarbonate  is  formed :  — 

CaCO3  +  H2O  +  CO2  =  Ca(HCO3)2. 

On  boiling  the  solution,  however,  carbon  dioxide  is  expelled 
and  calcium  carbonate  is  reprecipitated  :  — 

Ca(HCO3)2  =  CaCO3  +  H2O  4-  CO2. 

The  fact  that  limewater,  or  baryta  water  Ba(OH)2,  is  rendered 
turbid  by  carbon  dioxide  is  commonly  used  as  a  means  for 
detecting  the  latter,  though  in  some  cases  such  a  test  requires 
further  confirmation.  Breathing  into  limewater  renders  the 


FIG.  86. 


220  OUTLINES   OF  CHEMISTRY 

latter  turbid,  due  to  the  formation  of  a  precipitate  of  calcium 
carbonate,  and  this  demonstrates  the  presence  of  carbon  diox- 
ide in  the  breath. 

Many  natural  waters  contain  lime  and  magnesia  in  solution 
in  form  of  the  bicarbonates,  Ca(HCO3)2  and  Mg(HCO3)2. 
When  such  waters  are  boiled,  carbon  dioxide  escapes  and 
the  normal  carbonates,  CaCO3  and  MgCO3,  are  precipitated. 
Water  containing  carbonate  or  bicarbonate  of  calcium  in  solu- 
tion is  called  hard  water.  Since  boiling  decomposes  the  bicar- 
bonate and  thus  removes  some  of  the  calcium  salts  in  form  of 
a  precipitate  of  CaCO3,  it  is  clear  that  after  boiling  the  water 
has  lost  some  of  its  "hardness."  Thus  it  is  common  to  speak 
of  temporary  hardness  of  water,  which  can  be  removed  by  boil- 
ing, and  permanent  hardness,  which  remains  even  after  boiling. 
The  carbonates  of  the  alkalies,  like  sodium  carbonate  Na2CO3, 
and  potassium  carbonate  K2CO3,  may  be  fused  without  de- 
composition. But  the  carbonates  of  other  metals  are  com- 
monly decomposed  on  ignition,  yielding ,  carbon  dioxide  and 
the  oxide  of  the  metal,  thus :  — 

CaCO3  =  CaO  +  CO2. 
SrC03  =  SrO  +  CO2. 

Properties  of  Carbon  Dioxide.  —  Carbon  dioxide  is  a  colorless 
gas,  of  slightly  acid  taste  and  a  feeble,  agreeably  pungent  smell. 
It  is  1.529  times  as  heavy  as  air.  Being  so  heavy,  it  may 
be  readily  poured  or  even  siphoned  from  one  jar  to  another 
(Fig.  87).  It  neither  supports  combustion  nor  respiration.  At 
and  below  31°,  its  critical  temperature,  carbon  dioxide  may  be 
liquefied  by  pressure.  The  liquid  boils  at  —  79°  under  atmos- 
pheric pressure.  Solid  carbon  dioxide  readily  forms  when  the 
liquid  is  allowed  to  evaporate  rapidly  in  the  air,  for  thus  much 
heat  is  absorbed  and  the  remaining  liquid  is  chilled  below  the 
melting  point.  The  solid  looks  like  snow,  and  melts  at  —  57°. 
In  the  air  it  evaporates  without  first  passing  into  the  liquid 
state,  which  is  quite  natural,  since  its  boiling  point  under  at- 
mospheric pressure  is  much  lower  than  its  melting  point.  The 
large  amounts  of  carbon  dioxide  evolved  by  fermentation  in  the 
brewing  industries  are  now  collected,  washed,  and  pumped  into 
cylinders  made  of  mild  steel  (Fig.  13.)  In  this  process  the  car- 
bon dioxide  is  chilled  and  put  into  the  cylinders  under  pressure. 


CARBON  AND   SOME   OF   ITS  TYPICAL   COMPOUNDS        221 


It  is  used  for  making  soda  water,  which  consists  of  water  charged 

with  carbon  dioxide  under  pressure.     When  the  pressure  is  re- 

leased, a  portion  of  the  gas  escapes, 

for   under   atmospheric    pressure 

and  room  temperature,  water  dis- 

solves only  about  its  own  volume 

of  the  gas.      The  solution  has  a 

faintly  acid   reaction  toward  lit- 

mus.    The   name    "soda  water" 

comes  from  the  fact  that  sodium 

bicarbonate  NaHCO3,  also  popu- 

larly called  bicarbonate  of  soda  or 

simply  soda,  is  at  times  used  in 

preparing  carbonated  water. 

Solid  carbon  dioxide  is  fre- 
quently used  for  securing  low 
temperatures.  For  this  purpose 
it  is  often  employed  mixed  with 
ether  in  order  to  secure  more 
rapid  evaporation.  Thus  tem- 
peratures as  low  as  —  80°  may  be 
secured,  and  in  a  partial  vacuum 
even  —  100°  may  be  reached.  Natural  carbon  dioxide,  as  it 
issues  from  the  earth,  is  also  frequently  bottled  in  steel  cylin- 
ders as  described,  and  placed  on  the  market.  Thousands  of 
tons  of  carbon  dioxide  are  thus  sold  annually. 

Carbon  dioxide  is  also  used  as  a  fire  extinguisher.  The  reason 
why  it  does  not  burn  or  support  'combustion  is  that  it  contains 
all  the  oxygen  it  can  hold,  and  retains  this  very  tenaciously.  At 
high  temperatures  potassium  will  burn  brilliantly  in  carbon 
dioxide,  for  potassium  is  under  these  conditions  able  to  rob 
carbon  dioxide  of  its  oxygen,  thus  :  — 


FIG.  87. 


Here  again  we  see  the  relative  character  of  the  process  of 
combustion. 

Physiological  Effects  of  Carbon  Dioxide.  —  In  pure  carbon 
dioxide,  living  beings  will  die  for  lack  of  oxygen,  just  as  they 
would  succumb  in  nitrogen,  for  example.  Air  containing  over 
20  per  cent  of  carbon  dioxide  may  also  finally  produce  death 


222  OUTLINES   OF   CHEMISTRY 

when  breathed ;  for  in  respiration  carbon  dioxide  is  given  off 
and  oxygen  is  taken  up  from  the  air,  which  processes  are 
greatly  impeded  by  a  high  content  of  carbon  dioxide  in  the  air. 
Upon  the  mucous  membranes,  carbon  dioxide  has  a  stimulating 
action  which  is  agreeable  and  refreshing,  for  this  reason  car- 
bonated drinks  are  highly  esteemed.  Water  from  which  all 
carbon  dioxide  has  been  expelled  by  recent  boiling  tastes  flat, 
as  already  stated.  Carbon  dioxide  colors  the  blood  dark  brown, 
and  so  when  air  containing  even  but  6  per  cent  of  carbon  diox- 
ide is  breathed  continuously  drowsiness  results.  The  anaes- 
thetic effects  of  carbon  dioxide  are  similar  to  those  of  nitrous 
oxide.  These  effects  may  be  counteracted  by  taking  the  person 
affected  into  fresh  air  that  contains  only  the  usual  amount  of 
carbon  dioxide,  for  thus  the  blood  again  is  able  to  get  its 
needed  oxygen,  and  so  the  intoxication  gradually  disappears. 
The  necessity  of  good  ventilation  of  buildings,  especially  sleep- 
ing rooms  and  audience  halls,  is  consequently  apparent.  The 
evil  effects  of  re-breathing  carbon  dioxide  are  great  in  them- 
selves, but  the  exhalations  aside  from  their  carbon  dioxide  con- 
tent are  much  more  injurious  to  health. 

Relations  of  Carbon  Dioxide  to  Plant  and  Animal  Life.  —  All 
animals  exhale  carbon  dioxide,  which  is  produced  during  the 
life  process  as  a  result  of  slow  oxidation  of  their  tissues.  The 
heat  of  the  animal  body  is  a  result  of  the  oxidation  that  is  con- 
tinually going  on  while  the  animal  lives.  Now  the  carbon  diox- 
ide thus  exhaled  enters  the  atmosphere,  and  is  again  taken  up 
by  the  green  leaves  of  plants  in  which  in  the  sunlight  in  presence 
of  the  chlorophyll,  water  and  carbon  dioxide  react  with  each 
other,  forming  compounds  containing  carbon,  hydrogen,  and 
oxygen,  notably  starch  (CgH^Og)^,  and  free  oxygen,  the  latter 
being  exhaled.  Thus,  carbon  dioxide  is  decomposed  and  starch 
is  formed,  which  may  again  serve  as  food  for  animals,  in  whose 
bodies  it  is  oxidized  to  carbon  dioxide  and  water.  And  so  the 
cycle,  often  spoken  of  as  the  carbon  cycle,  repeats  itself  over 
and  over.  The  energy  is  furnished  by  the  sun's  rays,  which  pro- 
duce the  decomposition  of  carbon  dioxide  and  water  into  starch 
and  oxygen  in  the  green  leaf  of  the  plant.  The  starch  and  free 
oxygen  formed  contain  more  energy  than  the  carbon  dioxide 
and  water ;  and  this  excess  of  energy  in  the  starch  and  oxygen 
is  again  given  off  in  the  animal  body,  when  carbon  dioxide  and 


CARBON  AND   SOME   OF  ITS  TYPICAL  COMPOUNDS        223 

water  are  formed  as  a  result  of  the  oxidation  of  the  starch.  In 
connection  with  the  study  of  the  element  nitrogen,  we  have 
seen  that  here  too  a  somewhat  similar  cycle  occurs. 

Early  Work  on  Carbon  Dioxide.  —  J.  B.  van  Helmont  (1577- 
1644)  showed  that  the  same  gas  is  formed  during  alcoholic 
fermentation,  during  the  action  of  acids  on  chalk,  and  during 
the  combustion  of  carbon.  He  demonstrated  that  this  gas  is 
also  found  issuing  from  fissures  in  volcanic  regions,  and  that  it 
is  contained  in  the  waters  of  mineral  springs.  Indeed,  it  was 
he  who  first  used  the  term  gas.  He  called  carbon  dioxide  "  gas 
sylvestre  "  or  "  gas  carbonum."  He  was  aware  of  the  fact  that 
carbon  dioxide  does  not  support  combustion  or  respiration. 
Stephen  Hales  was  the  first  to  collect  gases  over  water  by  dis- 
placement, as  is  still  in  vogue,  while  Priestley  taught  how  to 
use  mercury  for  this  purpose  instead  of  water.  The  fact  that 
carbon  dioxide  is  absorbed  by  caustic  alkalies  was  discovered 
by  Joseph  Black  in  1757.  He  called  the  gas  "  fixed  air  "  be- 
cause it  is  thus  absorbed  or  fixed  by  caustic  alkalies.  Black 
demonstrated  that  soluble  salts  are  formed  when  carbon  dioxide 
acts  on  caustic  potash  or  soda,  and  that  an  insoluble  precipitate 
results  when  the  gas  is  conducted  into  limewater.  He  also 
found  that  carbon  dioxide  is  liberated  when  limestone  is 
strongly  heated,  as  in  the  process  of  making  lime.  McBride 
showed  that  the  gas  is  liberated  during  putrefaction.  Priestley 
demonstrated  its  presence  in  the  air,  and  Lavoisier  proved  that 
carbon  dioxide  is  formed  during  respiration  and  the  reduction 
of  metallic  oxides  by  charcoal.  The  latter  also  showed  that 
the  gas  contains  only  carbon  and  oxygen,  while  to  the  work  of 
Berzelius,  Dumas,  Stas,  and  Roscoe  we  owe  the  careful  deter- 
mination of  the  percentage  composition  of  carbon  dioxide. 

Carbon  Monoxide.  —  Carbon  monoxide  occurs  in  gases  that 
issue  from  volcanoes.  It  is  a  constituent  of  illuminating  gas 
and  particularly  of  so-called  water  gas.  Furthermore,  it  often 
occurs  in  the  gases  issuing  from  blast  furnaces  in  which  iron 
ores  or  other  metallic  oxides  are  being  reduced. 

The  gas  may  readily  be  formed  by  passing  carbon  dioxide 
over  red-hot  carbon,  thus  :  — 


The  reaction  is  reversible.     At  1000°  it  is  nearly  complete  in 


224  OUTLINES   OF   CHEMISTRY 

the  sense  of  the  upper  arrow,  while  at  450°  it  is  practically 
completely  reversed.  Carbon  monoxide  is  commonly  formed 
in  ordinary  coal  fires,  where  the  carbon  dioxide  passes  upward 
through  red-hot  layers  of  coal.  The  blue  flame  so  frequently 
observed  in  a  coal  fire  is  due  to  the  combustion  of  carbon  mon- 
oxide. This  gas  is  always  formed  to  some  extent  when  carbon 
is  burned  in  an  insufficient  supply  of  oxygen.  In  this  case, 
however,  it  is  never  pure,  being  always  associated  with  carbon 
dioxide.  The  latter  may  be  readily  removed  by  passing  the 
gases  through  caustic  potash  solution,  which  absorbs  the  car- 
bon dioxide,  but  not  the  carbon  monoxide. 

When  air  is  passed  through  beds  of  incandescent  coke  or 
coal,  in  furnaces  of  special  type,  the  issuing  gas  consists  of 
28  to  30  per  cent  carbon  monoxide,  63  per  cent  nitrogen,  and 
smaller  amounts  of  carbon  dioxide.  This  gas  is  known  as 
producer  gas  and  is  extensively  used  for  fuel,  being  readily 
obtained.  The  nitrogen  and  carbon  dioxide  greatly  dilute  the 
gas  and  diminish  its  heating  power. 

When  steam  is  passed  over  carbon  heated  to  1000°  to  1400°  C. 
(see  apparatus,  Fig.  4),  carbon  monoxide  and  hydrogen,  a  mix- 
ture which  is  known  as  water  gas,  is  produced,  thus  :  — 

C  +  H20  =  CO  +  H2. 

This  water  gas  is  frequently  made  on  a  large  scale  and  used  as 
a  fuel.  In  America  it  is  also  often  used  for  illuminating  pur- 
poses, in  which  case,  since  it  burns  with  a  non-luminous  flame, 
it  'must  be  enriched  or  "carbureted"  by  the  addition  of  the 
vapors  of  hydrocarbons  that  are  rich  in  carbon  (see  illuminat- 
ing gas). 

Carbon  monoxide  is  produced  when  many  metallic  oxides  are 
heated  with  excess  of  carbon,  thus  :  — 

ZnO  +  C  =  Zn  +  CO. 

Indeed,  it  was  by  means  of  this  reaction  that  carbon  monoxide 
was  first  observed  (1776). 

Carbon  monoxide  is  further  formed  by  heating  carbonates 
with  carbon  or  zinc  dust,  thus  :  — 


CaC03+C  =  CaO 
MgCO3  +  Zn  =  ZnO  +  MgO  +  CO. 


CARBON  AND   SOME   OF  ITS  TYPICAL  COMPOUNDS        225 

The  gas  is  also  formed  by  heating  many  organic  acids  with 
concentrated  sulphuric  acid,  which  abstracts  water  from  the 
organic  acids  and  so  decomposes  them.  The  resulting  carbon 
monoxide  generally  contains  carbon  dioxide,  which  may,  how- 
ever, be  removed  by  means  of  caustic  potash,  as  already  de- 
scribed. In  the  laboratory  carbon  monoxide  is  often  prepared 
by  heating  oxalic  acid  with  sulphuric  acid,  thus  :  — 

(COOH)2  +  H2S04  =  H2S04.H20  +  CO2  +  CO. 
When  formic  acid  is  used,  carbon  monoxide  only  is  obtained :  — 
"  HCOOH  +  H2SO4  =  H2S04.H2O  +  CO. 

Frequently,  when  larger  quantities  of  carbon  monoxide  are 
required  for  experimental  work,  the  gas  is  prepared  by  heating 
together  potassium  ferrocyanide  with  ten  times  its  weight  of 
concentrated  sulphuric  acid  in  a  flask  of  relatively  large 
capacity.  The  reaction  which  occurs  is  as  follows  :  — 

6  H2S04  +  K4Fe(CN)6  +  6  H2O 

=  2  K2S04  +  3  (NH4)2S04  +  FeSO4  +  6  CO. 

Properties  of  Carbon  Monoxide.  —  Carbon  monoxide  is  a 
colorless,  odorless,  tasteless  gas  which  is  0.9672  time  as 
heavy  as  air.  At  and  below  —  139.5°  it  may  be  condensed 
to  a  liquid  by  pressure.  The  critical  pressure  is  35.5  atmos- 
pheres. The  liquid  boils  at  —  190°,  and  the  white,  snowlike 
solid  melts  at  —  207°.  In  the  air  and  in  oxygen,  the  gas  burns 
with  a  rather  small  blue  flame,  forming  carbon  dioxide,  thus  :  — 

2CO  +  O2=2C02. 

(2vols.)     (1vol.)      (2vols.) 

It  is  found  by  exploding  a  mixture  of  carbon  monoxide  and 
oxygen  that  2  volumes  of  the  former  and  1  volume  of  the 
latter  form  2  volumes  of  carbon  dioxide,  as  indicated  in  the 
above  equation.  From  this  fact  and  the  one  that  22.4  liters  of 
carbon  monoxide  under  standard  conditions  weigh  28  grams,  it 
follows  that  the  formula  for  the  gas  is  CO.  The  composition 
of  carbon  dioxide  must,  of  course,  be  known  as  a  result  of 
independent  experiment.  In  the  volumetric  relation  of  the 
combination  of  carbon  monoxide  and  oxygen,  we  have  another 
excellent  illustration  of  the  law  of  Gay-Lussac  of  combination 
of  gases  by  volume,  which  law  serves  as  the  main  support  of 
Avogadro's  hypothesis. 


226  OUTLINES  OF   CHEMISTRY 

Carbon  monoxide  is  a  strong  reducing  agent.  It  is  able  to 
abstract  oxygen  from  many  metallic  oxides  at  higher  tempera- 
tures, thus : — 

CuO  +  CO  =  Cu  +  C02. 
Fe2O3  +  3  CO  =  2  Fe  4-  3  CO2. 

In  sunlight  carbon  monoxide  unites  with  chlorine,  forming  an 
addition  product,  phosgene  COC12;  while  on  heating  sul- 
phur vapor  with  carbon  monoxide,  carbon  oxysulphide  COS 
is  formed.  Phosgene,  or  carbonyl  chloride,  boils  at  -f  8°  and 
is  readily  decomposed  by  water :  — 

COC12  +  H2O  =  CO2  +  2  HC1. 

Carbon  oxysulphide  is  a  colorless,  inflammable  gas  with  an 
odor  like  that  of  hydrogen  sulphide.  When  burned,  the  prod- 
ucts formed  are  carbon  dioxide  and  sulphur  dioxide :  — 

COS  +  3O=CO2  +  SO2. 

With  nickel  and  iron,  carbon  monoxide  forms  the  carbonyl 
compounds  Ni(CO)4  and  Fe(CO)5. 

The  chemical  behavior  of  carbon  monoxide  is  readily  ex- 
plained by  regarding  it  as  an  unsaturated  compound,  the  two 
free  bonds  of  the  carbon  atom  enabling  the  formation  of  the 
various  addition  products  to  take  place.  On  the  other  hand, 
it  must  be  remembered  that  while  carbon  monoxide  is  a  strong 
reducing  agent,  it  may  itself  in  turn  be  reduced  by  still  stronger 
reducing  agents  like  magnesium  and  aluminum,  whose  oxides 
at  the  high  temperatures  at  which  the  reaction  takes  place  are 
very  stable  and  non- volatile,  thus  :  — 

3  CO  +  2  Al  =  A12O3  +  3  C. 
CO  +  Mg  =  MgO  +  C. 

In  these  reactions,  then,  carbon  monoxide  is  compelled  to  play 
the  role  of  an  oxidizing  agent,  which  again  shows  us  that  oxida- 
tion and  reduction  are  processes  that  are  relative  in  character. 
Carbon  monoxide  is  readily  absorbed  by  an  ammoniacal  or 
hydrochloric  acid  solution  of  cuprous  chloride  at  room  tempera- 
tures. The  latter  solution  is  much  used  in  estimating  carbon 
monoxide  in  gas  analysis.  From  these  solutions  carbon  mo- 
noxide may  be  expelled  by  heating. 


CARBON  AND  SOME   OF  ITS  TYPICAL  COMPOUNDS        227 

Physiological  Effects  of  Carbon  Monoxide.  —  Carbon  monoxide 
is  a  very  poisonous  gas,  and  is  all  the  more  dangerous  because 
it  is  odorless,  and  so  does  not  betray  its  presence  till  it  has 
already  produced  toxic  effects.  The  gas  unites  with  the  hemo- 
globin of  the  blood,  forming  an  addition  product  which  is  bright 
red  in  color  and  very  stable.  This  fact  was  discovered  in  1826 
by  Piorry.  The  carbon  monoxide  hemoglobin  is  much  more 
stable  than  the  oxy hemoglobin.  Air  containing  as  low  as  one- 
twentieth  of  one  per  cent  carbon  monoxide  exerts  toxic  effects. 
These  manifest  themselves  as  headache,  unconsciousness,  con- 
vulsions, and  finally  death,  which  is  caused  by  about  100  cc. 
of  pure  carbon  monoxide  for  every  10  kilograms  of  weight  of 
the  person  or  animal  inhaling  the  substance.  The  resuscitation 
of  persons  poisoned  by  inhaling  carbon  monoxide  is  effected  by 
means  of  fresh  air,  in  very  mild  cases.  In  serious  cases,  oxy- 
gen must  be  supplied,  preferably  under  pressure  of  from  one 
and  one  half  to  two  atmospheres.  As  carbon  monoxide  is 
commonly  produced  in  coal  stoves,  it  is  necessary  to  provide 
suitable  draught  to  burn  the  gas  completely,  or  at  any  rate  to 
carry  off  the  products  of  combustion  so  that  they  cannot  escape 
into  the  room.  Water  gas,  which,  as  we  have  seen,  contains 
about  50  per  cent  carbon  monoxide,  is  doubtless  more  poisonous 
than  ordinary  illuminating  gas  made  by  heating  coal. 

Carbon  Bisulphide.  —  When  sulphur  vapor  is  passed  over 
charcoal  or  coke  heated  to  redness,  carbon  bisulphide  CS2  is 
formed.  It  is  also  called  carbon  disulphide.  The  carbon  is 
heated  in  a  tall  iron  cylinder,  and  the  sulphur  vapors  are  passed 
upward  through  the  hot  coal,  the  product  being  conducted  off 
in  tubes  and  condensed  to  a  liquid.  The  heating  is  now  gen- 
erally accomplished  by  means  of  electricity.  The  continuous 
process  thus  devised  by  E.  R.  Taylor  represents  a  very  great 
improvement  over  older  methods.  Figure  88  represents  the 
Taylor  furnace.  Pieces  of  coke  are  placed  between  the  elec- 
trodes E,  which  are  supplied  with  a  strong  alternating  current. 
The  heat  produced  melts  and  volatilizes  the  sulphur  S,  which 
is  continuously  supplied  through  B.  The  coke  is  renewed 
through  O.  The  tower  is  filled  with  charcoal  which  is  intro- 
duced from  above  through  D.  The  carbon  disulphide  vapors 
pass  off  through  A  to  the  condensers.  This  furnace  is  40  ft. 
high  and  16  ft.  in  diameter. 


228 


OUTLINES   OF   CHEMISTRY 


Carbon  bisulphide  forms  a  colorless,  volatile  liquid  of  not  un- 
pleasant ethereal  odor  when  pure.     As  it  usually  comes  in  the 

market,  it  is  slightly 
yellowish  in  color 
and  has  a  disagree- 
able odor,  which  is 
due  to  impurities. 
These  readily  form 
on  standing,  espe- 
cially in  presence  of 
moisture,  because 
of  slight  decompo- 
sition of  the  carbon 
disulphide.  The 
latter  has  a  specific 
gravity  of  1.262  at 
20°;  its  vapor  is 
2.68  times  as  heavy 
as  air.  Carbon  di- 
sulphide is  an  ex- 
tremely inflammable 
liquid  and  is  con- 
sequently dangerous 
to  handle.  The 
vapors  catch  fire  in 
the  air  when  heated 
to  but  232°.  Mixed 
with  air,  its  vapors 
are  explosive.  Car 
boil  disulphide 
burns  with  a  blue 
flame,  forming  car- 
bon dioxide  and 
sulphur  dioxide :  — 


B 


FIG. 


Carbon  disulphide  is  a  good  solvent  for  fats,  oils,  iodine,  rubber, 
and  sulphur.  It  is  used  as  a  solvent  for  fats  and  oils  on  a 
large  scale,  also  for  vulcanizing  rubber.  It  is  further  employed 
in  exterminating  ants,  lice,  and  other  insect  pests.  When  in- 


CARBON  AND   SOME   OF   ITS  TYPICAL  COMPOUNDS        229 

haled,  its  vapors  act  as  an  anaesthetic,  large  quantities  producing 
intoxications  and  serious  disturbances  of  the  nervous  system. 

Carbon  disulphide  CS2  is  analogous  to  carbon  dioxide  CO2. 
Just  as  CO2  forms  carbonates  with  oxides  of  alkalies,  so  CS2 
forms  analogous  compounds,  namely  trithiocarbonates,  with  sul- 
phides of  alkalies  :  — 

CaO+CO2=CaCO3. 
CaS  +  CS2  =  CaCS3. 
K2S  +  CS2=K2CS3. 

By  treating  a  trithiocarbonate  with  a  dilute  acid,  trithiocarr 
bonic  acid  H2CS3  is  liberated  as  an  unstable  oil  which  decom- 
poses readily  into  CS2  and  H2S,  thus  showing  great  analogy  to 
H2C03:  — 

H2C03=H20 


H2CS3  =  H2S  -f-  CS2« 

Cyanogen.  —  Under  ordinary  conditions,  carbon  and  nitrogen 
do  not  combine  with  each  other  ;  but  at  high  temperatures  in 
presence  of  carbonates  of  the  alkalies  or  oxides  of  the  alkaline 
earth  metals,  cyanides  are  formed.  These  are  compounds  con- 
taining carbon,  nitrogen,  and  the  alkali  metal  employed.  Thus, 
nitrogen  passed  over  a  mixture  of  carbon  and  fused  potassium 
carbonate  yields  potassium  cyanide  :  — 

K2C03+  3  C  +  N2  =  2  KCN  +  CO  +  CO2. 

In  this  reaction,  metallic  potassium  is  formed,  which  then 
unites  with  the  carbon  and  nitrogen  to  form  potassium  cya- 
nide. Whenever  any  carbon  compound  containing  nitrogen 
is  heated  with  metallic  potassium,  potassium  cyanide  results,  which 
fact  is  used  as  a  test  for  nitrogen  in  organic  compounds.  When 
calcium  oxide  is  employed,  the  cyanide  of  calcium  is  formed  :  — 

CaO  +  3  C  +  N2  =  Ca(CN)2  +  CO. 

On  passing  ammonia  over  carbon  heated  to  redness,  ammonium 
cyanide  is  produced  :  — 

2NH3+C=NH4CN  +  H2. 

The  cyanides  are  salts  derived  from  hydrocyanic  acid  HCN 
(which  see).  On  heating  mercuric  C}^anide  Hg(CN)2,  it  decom- 
poses, yielding  mercury  and  cyanogen,  thus  :  — 

Hg(CN)2  =  Hg+(CN)2. 


230  OUTLINES   OF   CHEMISTRY 

This  reaction  is  quite  similar  to  that  of  the  decomposition  of 
mercuric  oxide  by  heat.  Cyanogen  is  an  extremely  poisonous, 
colorless  gas  of  sharp  characteristic  odor.  It  -may  be  condensed 
to  a  colorless  liquid  that  boils  a,t  —21°.  The  solid  melts  at 
—  35°.  The  gas  burns  with  a  beautiful  purple  flame.  In 
water,  the  gas  is  readily  soluble,  also  in  alcohol.  Cyanogen 
gets  its  name  from  the  fact  that  it  enters  into  a  number  of  com- 
pounds that  are  blue  in  color. 

Hydrocyanic  Acid.  —  Hydrocyanic  acid,  or  Prussic  acid,  has 
the  composition  HCN.  It  is  formed  when  potassium  cyanide  is 
treated  with  hydrochloric  acid  :  — 

KCN  +  HC1  =  KC1  +  HCN. 

It  may  also  be  prepared  by  treating  potassium  ferrocyanide 
K4Fe(CN)6  +  3  H2O  with  dilute  sulphuric  acid,  thus  :  — 

2  K4Fe(CN)6  +  3  H2SO4  =  6  HCN  +  3  K2SO4  +  K2Fe  •  Fe(CN)6. 

The  potassium  ferrocyanide  is  prepared  by  heating  animal 
refuse,  like  blood,  hoofs,  horns,  etc.,  with  iron  and  potassium 
carbonate.  The  product,  when  purified,  forms  beautiful  lemon- 
yellow  crystals  of  the  composition  K4Fe(CN)6-f  3  H2O.  It  is 
known  also  as  the  yellow  prussiate  of  potash. 

Hydrocyanic  acid  is  an  extremely  poisonous,  colorless,  mobile 
liquid,  which  smells  like  bitter  almonds.  It  boils  at  26°.  The 
colorless  crystalline  solid  melts  at  —11°.  Hydrocyanic  acid  is 
a  very  weak  acid  which  is  readily  soluble  in  water  and  alcohol. 
Its  salts  are  also  very  poisonous.  One  twentieth  of  a  gram  of 
hydrocyanic  acid  is  sufficient  to  cause  death  in  case  of  a  human 
being.  The  best  antidote  consists  of  a  3  per  cent  solution 
of  hydrogen  peroxide,  which  acts  thus :  — 

H202  +  2  HCN  =  H2NOC  -  CONH2. 

The  latter  compound  is  called  oxamide. 

Hydrocyanic  acid  is  used  to  kill  insects  that  infest  shrubs 
and  fruit  trees.  It  is  also  at  times  used  in  medicine.  It 
affects  chiefly  the  respiratory  organs. 

Cyanates  and  Sulphocyanates.  —  On  heating  potassium  cya- 
nide with  lead  oxide,  potassium  cyanate  KCNO  is  formed  :  — 

KCN  +  PbO  =*  Pb  +  KCNO. 


CARBON  AND   SOME   OF   ITS   TYPICAL   COMPOUNDS         231 

The  free  cyanic  acid  HCNO  is  a  liquid  which  readily  decom- 
poses into  ammonia  and  carbon  dioxide  when  treated  with 
water  :  — 

H2O  +  HCNO  =  NH3  +  CO2. 

Fused  with  sulphur,  potassium  cyanide  forms  potassium 
sulphocyanate  KCNS :  — 

KCN  +  S  =  KCNS. 

The  free  acid  HCNS  is  extremely  unstable.  With  ferric  salts 
potassium  sulphocyanate  forms  ferric  sulphocyanate  Fe(CNS)3, 
which  is  blood-red  in  solution,  and  serves  as  a  delicate  test  for 
ferric  salts. 

The  fact  that  cyanides  pass  over  into  cyanates  and  sul- 
phocyanates  shows  that  cyanides  are  really  unsaturated  in 
character. 


CHAPTER   XV 


HYDROCARBONS   AND   ADDITIONAL   COMPOUNDS   OF 

CARBON 

Hydrocarbons.  —  Hydrocarbons  are  compounds  of  hydrogen 
and  carbon.  They  are  very  numerous,  nearly  three  hundred 
being  known.  The  simplest  hydrocarbon  is  marsh  gas,  or 
methane  CH4.  It  issues  from  the  decaying  vegetable  matter 
in  ditches  and  marshes  on  warm  summer  days,  hence  its 
popular  name,  marsh  gas.  Methane,  together  with  hydrogen, 
is  a  prime  constituent  of  natural  gas,  which  is  found  in  the 
coal  and  oil  regions  of  Indiana,  Ohio,  Pennsylvania,  and  other 
states.  Methane  may  be  prepared  artificially  by  a  process  to  be 
described  later.  Petroleum  is  essentially  a  mixture  of  hydro- 
carbons. The  chief  petroleum  fields  are  located  in  Pennsyl- 
vania, New  York,  Ohio,  Indiana,  Kentucky,  Kansas,  Texas, 
Colorado,  California,  and  Canada.  Less  extensive  deposits 
are  found  in  Russia  near  the  Caspian  Sea  and  Black  Sea,  in 
China,  India,  and  Japan.  The  hydrocarbons  in  American 
petroleum  practically  all  belong  to  the  so-called  paraffin  series, 
of  which  methane  is  the  first  and  simplest  member.  The  fol- 
lowing table  gives  a  number  of  hydrocarbons  of  this  series  :  — 


NAME 

FORMULA 

BOILING  POINT 

CH, 

-  160° 

C9HR 

-  93° 

CoHo 

-  45° 

C,Hin 

+  1° 

37° 

69° 

Heptane  

98° 

Octane          .          

125° 

Nonane                   

149.5° 

It  will  be  noted  that  the  boiling  points  increase  with  the  car- 
bon content,  and  that  the    first  three  substances  are  gases  at 

232 


HYDROCARBONS  AND   THEIR   DERIVATIVES  233 

ordinary  temperatures.  The  difference  in  composition  between  any 
two  adjacent  members  is  CH2,  and  any  series  of  compounds  in 
which  this  is  the  case  is  called  an  homologous  series.  The  series 
above  given  is  the  so-called  normal  paraffin  series.  The  general 
formula  for  any  of  its  members  is  CnH2w  +  2.  Of  this  series, 
compounds  containing  up  to  sixty  carbon  atoms  are  known. 
The  higher  members  are  solids.  Thus  pentadecane  C15H32 
melts  at  +10°;  eicosane  C20H42  melts  at  36.7°;  heptacosane 
C27H56  melts  at  59.5°  ;  and  hexacontane  C60H122  melts  at  102°. 
In  the  members  of  this  series,  carbon  is  quadrivalent,  and  the 
compounds  are  called  saturated  hydrocarbons. 

When  crude  petroleum  is  placed  in  a  retort  and  subjected  to 
fractional  distillation  the  following  fractions  are  obtained  :  — 

Cymogene  (mainly  butane),  B.  P.  about      0°  C. 

Rhigolene  (butane  and  pentane),  B.  P.  about    16°. 
Petroleum  ether  (pentane  and  hexane),      B.  P.  about    50°  to    60°. 

Gasoline  (hexane  and  heptane),  B.  P.  about    70°  to    90°. 

Naphtha  (heptane  and  octane)  B.  P.  about    90°  to  120°. 

Benzine  (octane  and  nonane),  B.  P.  about  110°  to  140°. 

Kerosene  (nonane  to  heptadecane),  B.  P.  about  150°  to  300°. 

Naphtha  is  also  sometimes  called  ligroin.  Above  300°  heavy 
oils  pass  over  which  are  used  as  lubricating  oils.  At  still 
higher  temperatures,  vaseline  is  obtained,  which  consists 
mainly  of  C19H40  to  C21H44.  Finally,  from  the  residue  in  the 
retort  paraffin  is  separated  out  at  low  temperatures.  Paraffin 
is  a  white  waxlike  substance  consisting  of  C21H44  to  C32H66 
and  melting  from  45°  to  70°  C.  according  to  composition. 
Petroleum  ether  is  used  as  a  solvent,  also  in  making  illumi- 
nating gas,  and  as  an  anaesthetic.  For  the  latter  purpose  it 
must  be  specially  purified.  Sometimes  a  mixture  of  the  first 
four  products  named  in  the  last  table  is  called  petroleum  ether. 
Gasoline,  naphtha,  and  benzine  are  also  used  as  solvents  and 
frequently  as  fuels  in  stoves  and  engines.  They  are  also  em- 
ployed in  gas  manufacture.  Benzine  is  frequently  used  in 
paints  and  varnishes  as  a  substitute  for  turpentine.  Kerosene 
is  used  as  a  fuel  and  for  purposes  of  illumination.  There  are 
different  grades  of  kerosene.  These  vary  as  to  color  and  fire 
test  or  flash  point;  i.e.  the  temperature  at  which  evaporation 
is  sufficient  so  that  the  vapors  may  be  lighted  in  the  air.  Kero- 
sene having  a  flash  point  of  110°  F.  is  safe  for  use  in  lamps; 


234  OUTLINES  OF  CHEMISTRY 

this  is  the  standard  flash  test  fixed  by  law  in  most  of  the 
states,  but  sometimes  a  test  of  150°  F.  is  required.  In  purify- 
ing kerosene  it  is  washed  with  sulphuric  acid,  then  with  an 
alkali,  and  finally  with  water. 

Paraffin  lubricating  oils  are  now  very  much  used  particularly 
in  gas  and  gasoline  engines.  Vaseline  is  used  in  ointments  of 
various  kinds.  In  crude  form,  it  serves  as  cup  grease  and  axle 
grease.  Paraffin  is  used  in  making  candles  and  chewing  gum, 
also  in  making  paper  and  fabrics  waterproof,  in  insulating  wires 
and  electrical  apparatus  of  various  kinds,  in  manufacturing 
matches,  and  in  the  laundry.  Crude  oils  from  the  Texas  and 
California  fields  are  used  for  the  preservation  of  railroad  ties, 
which  are  saturated  with  the  oils  before  using. 

Hydrocarbons  may  be  prepared  in  several  ways.  The  com- 
mon method  of  making  methane  is  by  heating  sodium  acetate  with 
lime  or  caustic  soda :  — 

CHgCOONa  +  NaOH  =  Na2CO3  +  CH4. 

The  higher  hydrocarbons  may  be  prepared  in  a  similar  way  by 
using  salts  of  acids  of  higher  carbon  content.  Thus  to  prepare 
propane,  sodium  butyrate  is  heated  with  a  caustic  alkali :  — 

C3H7COONa  +  NaOH  =  Na2CO3  +  C3H8. 

Hydrocarbons  may  also  be  formed  by  treating  carbides  of  metals 
with  water.  Thus  aluminum  carbide  and  water  yield  meth- 
ane :  — 

A14C3  +  12  H2O  =  4  A1(OH)3  +  3  CH4. 

The  carbide  of  aluminum  may  be  prepared  by  heating  oxide  of 
aluminum  with  carbon  in  the  electric  furnace,  when  the  follow- 
ing reaction  occurs :  — 

2  A12O3  +  9  C  =  6  CO  +  A14C8. 

Other  simple  methods  of  preparing  hydrocarbons  consist  of 
treating  halogen  substitution  products  with  nascent  hydrogen 
or  with  sodium  :  — 

CH3I  +  2  H  =  HI  +  CH4. 

;  methyl  iodide  methane 

2  CH3I  +  2  Na  =  2  Nal  +  C2H6. 

ethane 

Hydrocarbons  act  neither  as  acids  nor  as  bases,  and  thus  they 
differ  materially  from  the  hydrides  of  such  elements  as  nitro- 
gen, sulphur,  and  the  halogens. 


HYDROCARBONS   AND   THEIR   DERIVATIVES  235 

Ethylene  C2H4,  also  called  olefiant  gas,  is  the  first  of  an 
homologous  series  of  unsaturated  hydrocarbons  whose  general 
formula  is  CnH2ra.  Ethylene  may  be  prepared  by  heating  alco- 
hol with  concentrated  sulphuric  acid,  which  abstracts  water 
from  the  alcohol  :  — 

C2H5OH  +  H2S04  =  H2S04.H20  +  C2H4. 

The  relation  of  ethylene  to  ethane  C2H6  is  readily  seen  from 
the  following  graphic  formulae  :  — 

H      H  H      H  H      H 

H-G-C-H,         C  =  C,       or       _  C  -  C  -. 
H      H  H      H  H      H 

ethane  ethylene 

When    ethane    is    treated    with    bromine,    substitution    takes 

place  :  — 

C2H6  +  Br2  =  C2H5Br  +  HBr  ; 

ethylene,  however,  reacts  as  follows  :  — 

C2H4  +  Br,  =  C2H4Br2. 

That  is,  ethylene  bromide  C2H4Br2,  a  colorless  liquid,  is 
formed,  showing  that  ethylene  is  unsaturated,  the  two  free 
bonds  manifesting  themselves  in  the  fact  that  the  hydrocarbon 
is  able  to  unite  with  two  atoms  of  bromine  by  simple  addition. 
This  is  characteristic  of  all  of  the  members  of  the  ethylene 
series,  and  is  often  expressed  by  the  so-called  double  bond, 

I        I 
—  C  =  C  —  ,  which,  however,   is   not  a  source   of   additional 

strength.  It  simply  indicates  that  the  compound  is  unsatu* 
rated,  i.e.  is  capable  of  forming  addition  products,  as  above 
illustrated.  Ethylene  gas  burns  with  a  luminous  flame. 

Acetylene  C2H2  is  the  lowest  member  of  an  homologous 
series  of  hydrocarbons  of  the  general  formula  CnH2n_2.  Acety- 
lene gas  may  be  produced  by  passing  ethylene  through  red-hot 
tubes  :  — 

C2H4  =  C2H2  +  H2. 

It  is  also  formed  to  some  extent  when  a  Bunsen  burner  burns 
below,  that  is,  has  "struck  back."  However,  the  gas  is  best 
prepared  by  the  action  of  calcium  carbide  on  water  :  — 


CaC2  +  2  H2O  =  Ca(OH)2  +  C2H 


2. 


236  OUTLINES   OF  CHEMISTRY 

Acetylene  burns  with  a  very  bright,  luminous  flame.  It  is  conse* 
quently  often  used  for  illuminating  purposes.  Acetylene  is 
still  less  saturated  than  ethylene,  which  fact  is  expressed  by 

I         I 
the  formula  HC  —  CH,  or  more  frequently  by  H  -  C  =  C  -  H. 

I         I 

The  triple  bond,  or  acetylene  bond,  indicates  the  unsaturated 
condition  of  the  compound,  i.e.  its  ability  to  unite  with  four 
additional  atoms  of  halogen,  for  instance. 

Benzene  C6H6  (not  to  be  confounded  with  the  petroleum 
product,  benzine,  which  is  entirely  different)  is  a  colorless, 
mobile  liquid  of  specific  gravity  0.8799  at  20°,  boiling  at  80°. 
The  compound  forms  crystals  that  melt  at  6°.  Benzene  is 
obtained  as  the  light  oil  from  coal  tar.  It  is  the  first  member 
of  an  homologous  series  of  hydrocarbons  of  the  general  formula 
CraH2ra_6.  It  burns  with  a  luminous  flame,  and  is  an  excellent 
solvent  for  fats,  resins,  iodine,  sulphur,  and  phosphorus.  From 
benzene  many  very  important  substances  like  carbolic  acid,  aniline, 
the  coal-tar  dyes,  and  many  medicinal  and  aromatic  substances 
are  derived.  Because  of  the  aromatic  odor  of  the  hydrocarbons 
of  the  benzene  series  and  their  derivatives,  they  are  commonly 
called  the  aromatic  series.  The  hydrocarbons  of  the  paraffin 
series  and  their  derivatives  are  also  known  as  the  fatty  series, 
for  the  fats  commonly  belong  to  this  series. 

Naphthalene  C10H8  forms  shining  leaflets  that  melt  at  79°. 
It  occurs  in  coal  tar  and  gives  the  latter  its  peculiar  odor. 
In  form  of  moth  balls  naphthalene  is  used  to  protect  woolen 
goods  from  moths.  Naphthalene  is  closely  related  to  benzene. 

General  Behavior  of  Hydrocarbons. — All  hydrocarbons  burn, 
and  when  the  combustion  is  complete,  the  products  formed  are  car- 
bon dioxide  and  water.  G-aseous  hydrocarbons,  or  the  vapors  of 
all  light  hydrocarbon  oils,  are  inflammable.  With  oxygen  or  air, 
they  form  mixtures  that  explode  when  brought  in  the  neigh- 
borhood of  a  flame.  The  larger  the  carbon  content  of  the 
molecule  of  a  hydrocarbon,  the  more  luminous  is  the  flame  with 
which  it  burns  (see  luminosity  of  flames). 

On  the  whole,  hydrocarbons  are  rather  inert  substances 
chemically.  They  are  practically  insoluble  in  water,  but  they 
are  miscible  with  one  another  in  all  proportions.  They  dissolve 
fats,  oils,  ether,  alcohol,  carbon  disulphide,  and  many  other  sub- 
stances of  kindred  character. 


HYDROCARBONS   AND  THEIR  DERIVATIVES  237 

Halogen  Substitution  Products. — When  methane  is  treated 
with  chlorine  in  the  sunlight,  the  hydrogen  atoms  may  be  re- 
placed one  after  another  by  chlorine  as  indicated  by  the  follow- 
ing equations :  — 

CH4     +  C12  =  CH3C1     +  HC1. 

methyl  chloride 

CH3C1  +  C12  =  CH2C12    +HC1. 

dichlormethane 

CH2C12  +  C12  =  CHC13     +  HC1. 

chloroform 

CHC13  +  C12  =  CC14        +  HC1. 

carbon  tetrachloride 

Chloroform,  a  heavy  colorless  liquid  boiling  at  61°,  is  very 
important  as  an  ansesthetic.  The  corresponding  bromine  com- 
pound, bromoform  CHBr3,  is  also  known.  It  is  a  liquid  boiling 
at  151°;  while  iodoform  CHT3,  a  yellow  crystalline  solid  melt- 
ing at  119°,  is  used  in  dressing  wounds. 

Carbon  tetrachloride  boils  at  76°.  It  is  not  inflammable  like 
gasoline,  though  like  the  latter  it  dissolves  fats  readily.  Hence 
carbon  tetrachloride  is  often  used  as  a  solvent,  particularly  for 
cleaning  clothes,  being  less  dangerous  to  handle  than  volatile 
hydrocarbons. 

We  have  already  seen  that  unsaturated  hydrocarbons  may 
form  simple  addition  products  with  halogens. 

Alcohols.  —  On  treating  methyl  chloride  with  caustic  potash, 
the  following  reaction  occurs  :  — 

CH3C1  +  KOH  =  KC1  +  CH3OH. 

methyl  alcohol 

The  radical  CH3  is  called  methyl.  There  are  many  similar 
hydrocarbon  radicals,  which  are  also  called  alkyl  radicals.  So, 
for  instance,  we  have  :  ethyl  C2H5,  in  ethyl  chloride  C2H5C1 ; 
propyl  C3H7,  in  propyl  iodide  C3H7I ;  phenyl  C6H5,  in  phenol 
CgH5OH.  These  radicals  may  pass  from  one  compound  to  an- 
other precisely  as  the  radical  ammonium  NH4  does  in  ammonium 
compounds. 

Now  C H3OH  or  methyl  hydroxide  is  methyl  alcohol.  It  is  also 
called  wood  alcohol,  or  spirit  of  wood,  for  it  may  be  obtained  as 
one  of  the  products  of  the  dry  distillation  of  wood.  We  note 
that  methyl  alcohol  is  an  hydroxide.  Indeed,  we  may  regard 
it  as  water  with  one  hydrogen  atom  replaced  b}'  methyl,  or  as 


238 


OUTLINES   OF   CHEMISTRY 


sodium  hydroxide  with  the  sodium  atom  replaced  by  methyl. 
All  alcohols  are  hydroxides  of  alkyl  radicals,  and  the  general  for- 
mula of  an  alcohol  is  R  —  OH  where  R  stands  for  an  alkyl  radical. 
Thus,  C2H5OH,  or  ethyl  alcohol,  is  ordinary  alcohol;  C3H7OH 
is  propyl  alcohol ;  C4H9OH,  butyl  alcohol;  and  so  on  up  the  homol- 
ogous series,  the  higher  members  of  which,  like  C16H33OH,  cetyl 
alcohol,  and  C30H61OH,  melissyl  alcohol,  are  waxlike  solids. 

Ethyl  alcohol,  or  ordinary  alcohol,  is  also  called  spirit  of  wine, 
for  it  may  be  obtained  by  distilling  wine.  When  thus  pre- 
pared it  contains  other  aromatic  substances  from  the  wine. 

Alcohol  is  contained  in  all 
fermented  liquors.  It  may 
readily  be  prepared  by  fer- 
mentation of  glucose  with 
yeast,  thus :  — 


C6HI206=2C2H6OH 


glucose 


alcohol 


FIG.  89. 


Figure  89  shows  a  common 
form  of  yeast  cells  as  they 
appear  under  the  microscope. 
Pure  alcohol  boils  at  78° 
and  solidifies  at  about  —  130°. 
Beers  contain  from  3  to  5  per 
cent  alcohol,  wines  from  8  to  20  per  cent,  and  brandy,  whisky, 
and  rum  from  45  to  65  per  cent. 

By  adding  wood  alcohol  or  other  poisonous  substances  to 
ethyl  alcohol,  the  latter  is  made  unfit  for  use  as  a  beverage,  and 
is  said  to  be  "denatured."  Usually  about  10  volumes  of  wood 
alcohol  and  half  a  volume  of  benzene  are  added  to  100  volumes 
of  90  per  cent  alcohol  to  make  so-called  denatured  alcohol.  The 
latter  may  be  used  as  fuel  or  for  purposes  of  manufacturing, 
without  the  payment  of  duty. 

When  ethyl  alcohol  is  treated  with  phosphorus  trichloride,  the 
following  change  takes  place  :  — 


3  C2H5OH 


PC13  =  P(OH)3 


3  C2H5C1. 


Thus  chlorine  is  substituted  for  the  OH  group.  All  alcohols 
undergo  a  similar  change  when  treated  with  either  phosphorus 
chloride,  bromide,  or  iodide.  The  halogen  takes  the  place  of  the 


HYDROCARBONS  AND   THEIR   DERIVATIVES  239 

hydroxyl,  and  phosphorous  acid  is  formed.  Indeed,  this  fact  is 
used  to  ascertain  whether  the  OH  group  is  present  in  a  com- 
pound. The  number  of  such  groups  may  also  be  determined 
thus  by  the  number  of  halogen  atoms  that  enter  the  molecule. 
Water  itself  reacts  perfectly  analogously  with  phosphorus 
trichloride,  as  is  evident  from  the  following  equations  :  — 

3  H2O  +  PC13  =  P(OH)3  +  3  HC1. 
3  C4H9OH  +  PC13  =  P(OH)3  +  3  C4H9C1. 

When  treated  with  sodium,  the  hydrogen  of  the  OH  group  of 
an  alcohol  is  replaced,  thus  : 

2  C2H5OH  +  2  Na  =  2  C2H5ONa  +  H2. 

sodium 
alcoholate 

This  is  perfectly  analogous  to  what  happens  when  water  is 
treated  with  sodium  :  — 

2  H20  +  2  Na  =  2  HONa  +  Ha. 

Just  as  we  have  hydroxides  of  the  metals  which  contain  more 
than  one  hydroxyl  group,  like  Ca(OH)2,  Bi(OH)3,  Sn(OH)4, 
so  we  also  have  alcohols  that  contain  two  or  more  hydroxyl 
groups.  Thus,  we  have  :  — 

CH2-OH 
I 

CH-OH 

CH2-OH  I 

CH2-OH  |  CH-OH 

glycol,  CH  —  OH  glycerine,  and  |  mannite. 

|  C 

CH2-OH  | 


CH2-OH  |  CH-OH 


CH-OH 

I 
CH2-OH 

Glycol  and  glycerine  are  rather  viscous  liquids  that  mix  with 
water  in  all  proportions.  They  are  slightly  sweet.  Moreover, 
the  sweet  taste  increases  as  we  go  up  the  series  from  glycol  to 
mannite.  Erythrite,  which  contains  four  carbon  atoms  and 
four  hydroxyl  groups,  and  arabite,  which  contains  five  carbon 
atoms  and  five  hydroxyl  groups,  are  also  well  known,  though 
they  are  of  no  practical  importance.  These  alcohols  are  called 
polyhydric  alcohols.  They  may  act  towards  acids  like  bases, 
forming  salts  which  are  called  esters  (which  see).  Mannite  is 


240  OUTLINES   OF   CHEMISTRY 

a  beautifully  crystalline  substance  which  readily  dissolves  in 
water,  but  not  in  hydrocarbons.  It  is  closely  allied  to  the  sugars 
(which  see).  Taken  internally,  mannite  acts  as  a  mild  purga- 
tive. All  the  polyhydric  alcohols  are  soluble  in  water,  and 
from  erythrite  up  they  are  solids  under  ordinary  conditions. 

Phenols.  —  The  hydroxyl  derivatives  of  benzene  C6H6  are 
called  phenols.  The  simplest  of  these  is  C6H5OH.  It  is  called 
phenol,  or  more  commonly  carbolic  acid.  It  is  really  not  an 
acid,  though  it  does  exhibit  acidic  properties  to  some  extent. 
Thus  with  caustic  alkalies  it  forms  phenolates : 

C6H5OH  +  KOH  =  C6H5OK  +  H2O. 

potassium 
phenolate 

Alcohols  do  not  form  alcoholates  when  treated  with  caustic 
alkalies.  It  will  be  recalled  that  it  is  necessary  to  treat  alcohols 
with  metallic  sodium  or  potassium  to  form  the  corresponding 
alcoholates. 

Carbolic  acid  crystallizes  in  long  needles  that  melt  at  42°  and 
turn  pink  when  exposed  to  the  air.  In  water  it  is  but  spar- 
ingly soluble,  about  1  part  dissolves  in  15  parts  of  cold  water ; 
but  in  alcohol  and  many  other  organic  liquids  it  dissolves  much 
more  copiously.  Carbolic  acid  has  a  characteristic  odor  and  is 
very  poisonous,  whence  its  use  as  a  disinfectant  and  antiseptic. 
When  brought  in  contact  with  the  skin  it  exerts  a  corrosive 
action. 

Among  the  phenols  containing  more  than  one  hydroxyl 
group  hydroquinone  C6H4(OH)2  and  pyrogallol  C6H3(OH)3, 
also  called  pyrogallic  acid,  are  of  importance  as  developers  in 
photography  (which  see).  Further,  in  the  incomplete  combus- 
tion and  dry  distillation  of  wood  there  are  formed,  along  with 
other  products,  phenols,  notably  guajacol  C6H4(OCH3)OH  and 
kreosol  C6H3(CH3)(OCH3)OH,  a  mixture  of  which  is  called 
creosote.  These  give  the  smoke  a  penetrating  odor  and  anti- 
septic value  that  is  used  in  preserving  meats,  sausages,  arnd 
fish. 

Aldehydes.  —  On  careful  oxidation  of  alcohols,  aldehydes  are 
formed,  which  are  compounds  containing  two  hydrogen  atoms 
less  than  the  alcohol  from  which  they  may  be  obtained,  whence 
the  name  alcohol  dehydrogenatum,  abbreviated  aldehyde.  So 
when  methyl  alcohol  is  partially  oxidized  by  means  of  potassium 


HYDROCARBONS  AND  THEIR   DERIVATIVES  241 

permanganate,  or  incomplete  combustion,  the  following  change 

takes  place  :  — • 

H 
I 

CH3OH  +  O  =  H20  +  C  =  O. 
I 
H 

The  product,  formaldehyde,  is  a  gas  which  may  be  condensed 
to  a  liquid  that  boils  at  —  21°.  It  has  a  penetrating,  suffocating 
odor,  acts  strongly  on  the  mucous  membranes,  and  is  a  powerful 
antiseptic.  In  water  it  dissolves  up  to  about  40  per  cent,  and 
it  is  this  aqueous  solution  that  is  sold  under  the  name  of  forma- 
line. It  is  used  as  a  disinfectant,  antiseptic,  or  preservative  in 
various  strengths  as  required.  All  the  aldehydes  are  very  active 
chemically,  being  specially  strong  as  reducing  agents.  We  may 
regard  formaldehyde  as  carbon  dioxide  with  one  oxygen  atom 
replaced  by  two  hydrogen  atoms  (see  formula  above).  By  some 
it  is  thought  that  formaldehyde  is  the  first  product  formed 
when  water  and  carbon  dioxide  act  upon  each  other  in  the 
green  leaf  of  the  plant  in  the  sunlight,  forming  starch  and 
liberating  oxygen. 

Just  as  the  hydrocarbons,  their  halogen  substitution  products, 
and  the  alcohols  form  homologous  series,  so  the  aldehydes  form 
similar  series.  Thus,  we  have  formaldehyde  HCHO,  acetic 
aldehyde  CH3CHO,  propionic  aldehyde  C2H5CHO,  etc. ;  further, 
benzole  aldehyde  C6H6CHO,  also  called  oil  of  bitter  almonds,  and 
its  homologues. 

When  the  hydrogen  atoms  of  the  methyl  group  in  acetic  alde- 
hyde are  replaced  by  chlorine,  CC13-CHO  is  formed.  This  is 
trichloraldehyde,  or  chloral.  It  readily  unites  with  water,  form- 
ing the  white  crystalline  addition  product  CC13CH(OH)2,  chloral 
hydrate,  which  is  so  much  used  in  medicine  as  a  soporific. 

The  general  formula  of  an  aldehyde  is    R  —  C  =  O,  in  which  R 

I 

H 
represents  either  hydrogen  or  an  alkyl  radical. 

Organic  Acids.  —  On  further  oxidation  of  aldehydes,  acids  are 
produced.  In  this  process  one  molecule  of  aldehyde  takes  up 
an  additional  atom  of  oxygen.  The  reactions  in  the  formation 
of  formic  and  acetic  acids  from  formic  and  acetic  aldehydes  may 
be  represented  as  follows  :  — 


242  OUTLINES   OF   CHEMISTRY 

H-C=O     +O=     H-C=O; 
I  I 

H  OH 

formic  aldehyde  formic  acid 

CH3  -  C  =  O     +  O  =     CH3  -  C  =  O. 

I  I 

H  0-H 

acetic  aldehyde  acetic  acid 

In  general,  the  reaction  may  be  represented  thus  :  — 

R_C  =  O     +O=     R-C  =  O, 
I  I 

H  OH 

aldehyde  organic  acid 

R  representing  either  hydrogen  or  any  organic  radical.  The 
last  formula  above  given  is  the  general  formula  of  an  organic  acid. 
The  characteristic  group  which  it  contains,  namely,  —  C  =  O 

I 
OH 

or  CO  OH,  is  called  the  carboxyl  group.  The  hydrogen  in  this 
group  is  replaceable  by  metals  or  radicals,  just  as,  for  instance, 
the  hydrogen  in  nitric  acid  may  be  thus  replaced. 

Only  a  few  typical  organic  acids  can  be  mentioned  here. 
Formic  acid  H-COOH  occurs  in  red  ants  and  stinging  nettles. 
It  is  a  colorless  liquid  boiling  at  101°.  It  is  soluble  in  water 
in  all  proportions,  has  a  pungent  odor,  and  blisters  the  skin. 
With  bases  it  forms  the  formates,  thus  :  — 

HCOOH  +  NaOH  =  HCOONa  +  H2O. 

sodium  formate 

The  acid  is  consequently  monobasic,  only  one  hydrogen  atom 
being  replaceable  by  a  metal.  When  heated  in  closed  vessels 
to  160°,  formic  acid  yields  carbon  dioxide  and  hydrogen  :  — 

HCOOH  =  C02  +  H2. 

Sodium  formate  may  be  obtained  by  passing  carbon  monoxide 
over  heated  caustic  soda,  and  from  the  formate  the  free  acid 
may  be  obtained  by  means  of  sulphuric  acid,  thus :  — 

NaOH  +  CO  =  HCOONa ; 
HCOONa  +  H2S04  =  NaHSO4  +  HCOOH. 

Acetic  acid  CH3COOH  occurs  in  combination  with  organic 
radicals  in  many  odoriferous  plant  oils.  It  is  formed  as  one  of 


HYDROCARBONS  AND   THEIR   DERIVATIVES 


243 


the  products  of  the  dry  distillation  of  wood.     It  is  made  on  a 

large  scale  in  vinegar  factories,  the  process  depending  on  (1) 

the  formation  of  alcohol  by 

fermentation    of    sugar    pro- 

duced   from    the    starch    in 

grain,  and  (2)  the  oxidation 

of  this  alcohol  to  acetic  acid, 

which   is   brought  about   by 

Mycoderma  aceti,  "mother  of 

vinegar,"  a  bacterium  shown 

in  Fig.  90.     In  practice,  the 

dilute   alcohol,  8   to   15   per 

cent,   is    allowed    to    trickle 

over  beech  wood  shavings  con- 

tained in  a  barrel.     Thus  the 

alcohol  is  thoroughly  exposed 

to  the  oxygen  of  the  air,  and  the  acetic-acid-forming  bacteria 

cause  the  oxidation  to  take  place.     The  oxidation  process  may 

be  represented  thus  :  — 


FIG.  90. 


C2H5OH 


O2  =  H2O 


CHCOOH. 


When  alcohol  is  treated  with  oxygen  alone,  this  process  does 
not  occur  ;  however,  the  spores  of  the  acetic-acid-forming  bac- 
teria are  commonly  present  in  the  air,  and  so  various  alcoholic 
solutions  like  beer  and  wines  get  sour  because  of  the  oxidation 
of  the  alcohol  to  acetic  acid.  Cider  slowly  ferments,  forming 
alcohol,  which  is  then  similarly  converted  into  acetic  acid. 
Vinegar  obtained  as  above  described  is  a  solution  containing 
essentially  4  to  10  per  cent  acetic  acid.  Anhydrous  acetic  acid* 
may  be  obtained  from  this  solution  by  forming  sodium  acetate 
and  then  treating  the  latter  salt  with  sulphuric  acid  and  dis- 
tilling. The  reactions  are  :  — 

Na2C03  =  H2O  +  CO2  +  2  CH3COONa. 


2  CH3COOH 

(in  solution) 

CHoCOONa 

(dry  salt) 


H2S04  =  NaHSO4  +  CH3COOH. 

(glacial  acetic  acid) 


The  process  is  then  quite  similar  to  the  preparation  of  hydro- 
chloric acid. 

Pure  acetic  acid  boils  at  118°.     Jt  forms  crystals  that  melt 
at  16.5°,  so  that  it  is  easy  to  cause  it  to  solidify  on  a  cold  day, 


244  OUTLINES   OF  CHEMISTRY 

whence  the  popular  name  glacial  acetic  acid.  The  acid  has  a 
pungent  odor  and  corrodes  the  skin.  It  is  a  monobasic  acid, 
only  one  hydrogen  atom  being  replaceable  by  a  basic  element 
or  radical.  The  salts  formed  are  called  acetates. 

Among  the  homologues  of  acetic  acid  are  propionic  acid 
C2H6COOH,  butyric  acid  C3H7COOH,  which  smells  like  rancid 
butter  (in  which  it  occurs),  palmitic  acid  C15H31COOH,  and 
stearic  acid  C17H35COOH.  The  last  two  are  solids  at  ordinary 
temperatures.  They  occur  in  fats  and  oils  of  plant  and  animal 
origin,  commonly  together  with  oleic  acid  C17H33COOH,  an 
unsaturated  acid  belonging  to  another  homologous  series. 
These  acids  occur  in  fats  and  oils  as  esters  (which  see),  not  as 
free  acids. 

Benzole  acid  C6H5COOH  is  the  first  of  an  homologous  series 
of  acids  derived  from  benzene.  It  occurs  in  gum  benzoin  and 
Peru  balsam ;  and  a  derivative  of  benzoic  acid,  hippuric  acid 
C6H5-CO-NH-CH2-  COOH,  occurs  in  the  urine  of  herbivorous 
animals.  Benzoic  acid  crystallizes  in  shining  flakes  that  melt 
at  121.5°.  The  boiling  point  is  249°,  but  the  acid  sublimes 
readily  even  at  100°.  In  alcohol,  ether,  and  hot  water  it  readily 
dissolves,  while  in  cold  water  it  is  but  sparingly  soluble,  1  part 
in  400.  Its  odor  is  very  characteristic,  and  its  vapor  when 
inhaled  irritates  the  throat  and  nasal  passages,  causing  coughing 
and  sneezing.  It  is  much  used  as  an  antiseptic  and  preserva- 
tive, commonly  in  the  form  of  sodium  benzoate  C6  H5COONa. 
The  free  acid  forms  readily  by  oxidation  of  benzaldehyde. 

Oxalic  acid  is  the  simplest  dibasic  organic  acid.  Its  formula 
is  H2C2O4,  or 

COOH 

I          ; 
COOH 

that  is,  it  consists  of  two  carboxyl  groups.  It  is  readily  formed 
by  oxidizing  sugar,  sawdust,  glycol,  fats,  and  many  other 
organic  substances  with  strong  nitric  acid.  On  a  commercial 
scale,  sawdust  is  employed.  The  normal  sodium  salt  may  be 
obtained  by  passing  carbon  dioxide  over  sodium  at  350°,  thus:  — 

2  CO2  +  2  Na  =  (COONa)2,  i.e.  Na2C2O4. 

This  salt  is  also  formed  when  sodium  formate  is  quickly  heated 
to  250°,  thus  :  — 

2  HCOONa  =  H2  +  Na2C2O4. 


HYDROCARBONS  AND  THEIR  DERIVATIVES  245 

The  calcium  salt  is  difficultly  soluble  in  water,  and  this  fact  is 
often  used  in  analytical  chemistry,  in  detecting  and  estimating 
calcium.  When  calcium  oxalate  is  heated,  there  is  first  formed 
calcium  carbonate  and  carbon  monoxide,  after  which  the  car- 
bonate decomposes  to  calcium  oxide  arid  carbon  dioxide  on 
further  heating,  thus  :  — 

CaC2O4  =  CaCO3  +  CO, 
CaCO3  =  CaO  +  CO2  ; 
or  CaC2O4  =  CaO  +  CO  +  CO2. 

This  behavior  is  typical  of  the  oxalates  of  many  other  metals. 

Oxalic  acid  is  oxidized  to  carbon  dioxide  and  water  by 
oxidizing  agents  like  potassium  permanganate,  for  example, 
thus  :  — 

COOH 


COOH 


=  2  CO2  +  H2O. 


It  is  evident,  then,  that  oxalic  acid  is  a  good  reducing  agent.  In 
the  market  the  acid  is  commonly  sold  in  form  of  its  white  crys- 
talline hydrate  H2C2O4-2H2O.  By  heating  this,  the  anhy- 
drous acid  may  be  obtained,  which  on  further  heating  passes 
over  into  carbon  dioxide  and  formic  acid  :  — 

(COOH)2  =  CO2  +  HCOOH. 

Just  as  there  are  homologous  series  of  monobasic  organic 
acids,  so  there  are  homologous  series  of  the  dibasic  acids. 
So  we  have  malonic  acid  CH2(COOH)2,  succinic  acid 
(CH2COOH)2,  etc. 

Hydroxycarboxylic  acids.  —  The  simplest  of  these  is  glycolic 
acid,  or  hydroxy  acetic  acid,  CH2O  HCOOH,  which  may  be  pro- 
duced by  careful  oxidation  of  glycol  :  — 

CH2OH  CH2OH 

I  +20=|  +  H2O. 

CH2OH  COOH 

glycol  glycolic  acid 

Glycolic  acid  is  monobasic,  only  the  hydrogen  atom  in  the 
carboxyl  group  being  replaceable  by  a  basic  element  or  radical. 
The  other  OH  group  is  alcoholic  in  character.  This  compound 
is  therefore  both  an  alcohol  and  an  acid.  Its  homologue, 
C2H4.OH-COOH,  is  lactic  acid.  It  is  of  great  practical 


246 


OUTLINES   OF   CHEMISTRY 


importance.  In  its  pure  form  it  is  a  thick,  colorless,  odorless, 
hygroscopic  liquid  of  pronounced  acidic  properties.  The  acid  is 

monobasic,  and  forms  the  lactates  with  bases,  thus  :  — 

• 

C2H4  •  OH  •  COOH  +  KOH  =  H2O  +  C2H4OHCOOK. 

Lactic  acid  is  readily  miscible  with  water,  alcohol,  and  ether 
in  all  proportions.  Like  other  hydroxyacids,  it  cannot  be  dis- 
tilled, for  it  decomposes  into  various  simpler  products  like  CO, 
H2O,  CH3CHO,  etc.,  on  heating.  Lactic  add  is  the  acid  that 
causes  the  acidity  of  sour  milk,  whence  the  name  lactic  acid. 
Lactic  acid  is  produced  by  a  special  form  of  fermentation 
caused  by  lactic-acid-forming  bacteria.  These  are  shown  in 

Fig.  91,  together  with  a  few 
yeast  cells  to  indicate  approxi- 
mately the  relative  size  of  the 
organisms.  By  the  action  of 
these  bacteria,  starch  and  sug- 
ars are  converted  into  lactic 
acid,  which  is  consequently 
formed  in  many  liquids  that 
contain  starch,  sugars,  or  kin- 
dred organic  substances.  So, 
for  example,  lactic  acid  occurs 
in  sauerkraut,  in  sour  pickles, 
in  fermented  beet  juice,  and 
at  times  in  the  contents  of  the 
alimentary  tract.  The  lactates  of  strontium  and  of  silver  are 
used  in  medicine. 

Lactic  acid  also  occurs  in  muscular  tissues,  but  this  acid  is 
not  quite  identical  with  that  in  sour  milk.  So  the  acid  in  the 
muscles,  sarcolactic  acid,  has  the  power  to  rotate  the  plane  of 
polarized  light ;  that  is,  it  possesses  optical  activity,  a  property 
not  exhibited  by  lactic  acid  produced  by  fermentation,  which 
is  consequently  termed  inactive  lactic  acid.  Sarcolactic  acid 
turns  the  plane  of  polarized  light  to  the  right.  When  the 
fungus  Penicillium  glaucum  is  allowed  to  grow  in  solutions  of 
fermentation  lactic  acid,  the  latter  also  acquires  the  power  to 
turn  the  plane  of  polarized  light  to  the  right,  i.e.  it  becomes 
dextroactive,  and  is  in  every  way  identical  with  sarcolactic 
acid.  Now  when  the  bacillus  Acidi  Icevolactici  feeds  upon 


FIG.  91. 


HfDROCARBONS   AND   THEIR  DERIVATIVES  247 

solutions  of  fermentation  lactic  acid,  the  latter  acquires  the 
power  to  turn  the  plane  of  polarized  light  to  the  left,  that  is, 
it  becomes  Icevoactive.  Fermentation  lactic  acid  is  conse- 
quently a  mixture  of  equal  parts  of  dextro  and  laevo  lactic 
acid,  which  accounts  for  its  optical  inactivity.  The  optical 
activity  is  produced  as  described,  because  one  organism  feeds 
on  the  dextro  variety  of  the  acid,  whereas  the  other  organism 
lives  upon  the  Isevo  variety.  The  formula  of  lactic  acid  is 

H 

I 

H-C-H 

I 
H  -  C  -  OH. 

I 

c  =  o 

I 

O  -  H 

This,  like  all  other  chemical  formulae,  has  been  established  by 
a  study  of  the  methods  of  synthesizing  the  compound,  and  by 
its  deportment  toward  various  reagents.  It  will  be  observed 
that  lactic  acid  has  one  carbon  atom,  the  center  one,  whose  four 
bonds  are  satisfied  by  four  different  elements  or  radicals.  Such 
a  carbon  atom  is  called  an  asymmetric  carbon  atom.  All  carbon 
compounds  that  are  optically  active  possess  at  least  one  asymmetric 
carbon  atom  in  the  molecule ;  though,  as  we  have  seen,  the  pos- 
session of  such  an  asymmetric  carbon  atom  does  not  necessarily 
make  the  compound  optically  active,  for  the  substance  under 
consideration  may  be  a  mixture  of  equal  parts  of  the  dextro 
and  laevo  varieties,  as  in  the  case  of  fermentation  lactic  acid. 
In  general,  for  every  dextroactive  compound  there  is  a  laevo- 
active  compound  that  rotates  the  plane  of  polarized  light  in 
the  opposite  direction  to  the  same  degree.  In  order  to  repre- 
sent the  difference  between  dextro  and  laevo  compounds  by 
means  of  formulae,  Le  Bel  and  van't  Hoff  simultaneously  and 
independently  of  each  other  (1874)  proposed  to  represent  the 
carbon  atom  by  a  regular  tetrahedron  the  corners  of  which  in- 
dicate the  four  valences.  Figure  92  shows  the  two  formulae  for 
dextro  and  laevo  lactic  acids. 

It  will  be  observed    that   the  formulae   are  alike,  but   not 
superposable ;  that  is,  they  are  to  each  other  as  the  right  hand 


248 


OUTLINES   OF  CHEMISTRY 


is  to  the  left  hand,  or  as  an  object  is  to  its  image  in  the  mirror. 
The  study  of  the  composition  of  carbon  compounds  by  the  aid 
of  formulae  thus  represented  in  three  dimensions  has  been 
carried  on  with  marked  success,  and  has  in  recent  years  also 
been  extended  to  the  investigation  of  compounds  of  other 
elements,  notably  those  of  nitrogen.  The  branch  of  chemistry 
which  seeks  to  further  inquiry  by  the  use  of  formulse  in  three 
dimensions  is  called  stereo-chemistry. 

Compounds  that  contain  the  same  elements  in  the  same  pro- 
portions by  weight  are  said  to  be  isomers.  So  acetic  acid 
CH3COOH  and  methyl  formate  HCOOCH3-are  isomers,  for 


COOH 


COOH 


FIQ.  92. 


they  contain  the  same  elements  in  the  same  proportions  by 
weight,  and  yet  they  are  quite  different  compounds.  Dextro 
and  Isevo  lactic  acids  also  contain  the  same  elements  in  the 
same  proportions  by  weight,  and  in  addition  they  are  in  every 
way  identical  in  their  chemical  behavior.  The  difference  be- 
tween them  lies  simply  in  their  behavior  toward  polarized  light 
and  they  are  consequently  called  optical  isomers,  or  stereo-iso- 
mers,  and  the  property  they  thus  exhibit  is  called  stereo-isom- 
erism. 

When  an  aqueous  solution  of  formaldehyde  HCHO  is  evapo- 
rated, an  amorphous  white  substance  (HCHO)2,  paraformalde- 
hyde,  separates  out.  On  careful  heating,  this  may  be  transformed 
into  metaformaldehyde  (HCHO)3,  a  crystalline  compound  melt- 
ing at  171°,  which  on  being  heated  to  140°  with  much  water  is 
again  decomposed  into  formaldehyde.  Other  aldehydes  exhibit 
similar  properties.  The  process  of  forming  larger  molecules 
by  simple  aggregation  of  two  or  more  molecules  is  called  poly- 
merization, and  the  new  products  formed  are  said  to  be  polymers. 
Polymerization  is  quite  a  common  process.  It  will  be  noted 
that  polymers  contain  the  same  elements  in  the  same  proper- 


HYDROCARBONS  AND  THEIR  DERIVATIVES 


249 


tions  by  weight  as  the  simple  compounds  from  which  they  have 
sprung;  they  are  consequently  isomers  as  above  defined.  When, 
however,  the  molecular  weight  of  two  isomeric  compounds  is 
identical,  they  are  said  to  be  metameric,  in  contradistinction  to 
polymeric,  which  latter  term  is  applied  when  the  isomerism 
depends  upon  difference  in  molecular  weight,  as  already  ex- 
plained. 

The  optical  activity  of  substances  is  examined  by  means  of  a 
polariscope,  also  called  a  polarimeter,  a  common  form  of  which 


FIG.  93. 

is  shown  in  Fig.  93.  The  substance  to  be  tested,  commonly  in 
the  form  of  a  liquid  or  solution,  is  placed  in  a  tube  between  the 
polarizer  and  analyzer;  that  is,  the  nicol  prisms  which  consti- 
tute the  vital  part  of  the  instrument.  The  degree  of  rotation 
is  read  off  on  the  scale  attached  to  the  analyzer,  the  yellow 
sodium  flame  being  usually  employed  as  a  source  of  monochro- 
matic light.  The  rotation  is  proportional  to  the  length  of  the 
tube  and  the  concentration  of  the  solution  of  other  liquid 
employed.  By  concentration  is  meant  the  number  of  grams 
of  active  substance  contained  in  one  cubic  centimeter.  The 


250  OUTLINES   OF   CHEMISTRY 

specific  rotatory  power  of  any  substance  is  the  number  of  degrees 
of  rotation  it  exhibits  in  a  tube  1  decimeter  long  when  1  gram  of 
active  substance  is  contained  in  1  cc.  Many  important  substances 
of  commerce,  notably  sugars,  essential  oils,  and  other  com- 
pounds obtained  from  plants,  possess  optical  activity.  The 
polariscope  serves  in  detecting  the  presence  of  such  substances, 
and  also  in  estimating  their  amounts  when  present.  So,  for 
example,  in  estimating  the  amount  of  sugar  in  the  juice  of 
sugar  beets,  the  polariscope  enables  one  to  obtain  very  rapid 
and  accurate  results. 

Malic,  tartaric,  and  citric  acids  are  important  fruit  acids. 
Malic  acid  occurs  in  sour  apples,  in  mountain  ash  berries,  and 
in  many  other  fruits.  It  is  monohydroxysuccinic  acid : 

H 

I 
HO-C-COOH 

I 
H-C-COOH 

I 
'   H 

From  the  formula  it  is  evident  that  the  acid  is  dibasic,  and 
that  its  molecule  contains  an  asymmetric  carbon  atom.  The 
acid  is  optically  active.  This  is  also  true  of  tartaric  acid, 
which  is  dihydroxysuccinic  acid :  — 

H 
1 
HO  -  C  -  COOH 

I 
HO  -  C  -  COOH 

I 
H 

This  acid  occurs  in  grapes  as  the  acid  potassium  salt  C4H5O6K, 
which  is  difficultly  soluble  in  water  and  is  commonly  known  as 
cream  of  tartar.  In  its  pure  form  this  salt  is  perfectly  white  ; 
but  in  its  crude  state  it  commonly  has  a  brownish  red  appear- 
ance from  the  coloring  matter  of  the  grapes.  In  this  crude 
state  it  is  called  argol.  Sodium  potassium  tartrate  C4H4O6NaK 
is  called  Rochelle  salt.  It  is  used  as  a  mild  purgative.  The 
acid  sodium  salt  and  the  normal  sodium  and  potassium  salts  of 
tartaric  acid  are  readily  soluble  in  water.  They  are  all  optic- 
ally active.  The  common  variety  of  tartaric  acid  is  dextro- 


HYDROCARBONS   AND   THEIR   DERIVATIVES 


251 


active.  It  forms  beautiful  monoclinic  prisms.  The  dextro  and 
Isevo  varieties  crystallize  in  forms  that  are  to  each  other  as  an 
object  is  to  its  im- 
age in  the  mirror; 
that  is,  they  are 
enantiomorphous 
forms.  The  fact 
that  dextro  and 
Ise  vo  tartaric  acids 
also  exhibit  dex- 


tro and  Isevo  char- 
acter     in      their 


FIG.  94. 


crystal  forms  was  discovered  by  Louis  Pasteur.  It  was  his 
first  notable  scientific  discovery.  Figure  94  shows  crystals  of 
dextro  and  Isevo  tartaric  acid. 

In  lemons  and  other  citrus  fruits  citric  acid  abounds.  This 
is  a  strong  tribasic  hydroxyacid  whose  composition  is  repre- 
sented by  the  following  formula  :  — 

H 
I 
H  -  C  -  COOH 

I 
HO-  C-COOH. 

I 

H  -  C  -  COOH 
I 
H 

It  crystallizes  with  one  molecule  of  water  in  beautiful  rhombic 
prisms.  These  melt  at  100°  and  lose  their  water  at  130°. 
This  acid* forms  three  series  of  salts,  for  there  are  three  re- 
placeable hydrogen  atoms  in  the  molecule.  Thus  we  have: 
C6H5O7K3,  the  normal  or  tertiary  potassium  citrate;  C6H6O7K2, 
secondary  potassium  citrate  ;  and  C6H7O7K,  primary  potassium 
citrate. 

Malic,  tartaric,  and  citric  acids  were  discovered  by  the 
great  Scheele.  He  prepared  them  by  treating  fruit  juices 
with  lime,  thus  obtaining  the  calcium  salts,  and  then  decom- 
posing these  with  sulphuric  acid,  forming  calcium  sulphate, 
which  is  difficultly  soluble,  and  the  free  acids.  The  latter 
remained  in  the  solutions,  from  which  they  were  obtained  by 
evaporation. 


252  OUTLINES   OF  CHEMISTRY 

Esters.  — Esters,  or  ethereal  salts,  are  formed  by  replacing  the 
hydrogen  of  an  acid  by  an  alkyl  radical.  Just  as  metallic  hy- 
droxides react  with  acids  forming  salts  and  water,  so  alcohols, 
which  are  hydroxides  of  alkyl  radicals,  react  with  acids  to  form 
esters  and  water:  — 

NaOH  +  HCOOH  =  HCOONa  +  H2O. 
CH3OH  +  HCOOH  =  HCOOCHg  +  H2O. 

The  compound  HCOONa  is  sodium  formate,  whereas  HCOOCH3 
is  methyl  formate.  The  latter  is  a  typical  ester.  In  making 
esters,  it  is  common  practice  to  add  a  dehydrating  agent  to  take 
up  the  water  formed;  usually  hydrochloric  acid  gas  is  em- 
ployed, for  it  has  great  affinity  for  water,  and  can  readily  be 
removed  afterwards  because  of  its  volatility.  Ethereal  salts 
of  either  inorganic  or  organic  acids  may  be  formed.  Thus, 
we  have  such  compounds  as  methyl  iodide  CH3T,  ethyl 
nitrate  C2H5NO3,  ethyl  nitrite  C2H5NO2,  methyl  hydrogen 
sulphate  CH3HSO4,  amyl  acetate  CH3COOC5Hn,  propyl  tartrate 
C4H4O6(C3H7)2,  etc.  Many  of  the  simpler  esters  are  mobile 
liquids  of  pleasant  aromatic  odor,  which  can  readily  be  dis- 
tilled and  purified.  The  general  formula  of  an  ester  of  an  organic 
acid  is  RCOOR',  where  R  and  R'  represent  alkyl  radicals  that 
may  or  may  not  be  alike.  Esters  occur  in  flowers,  fruits,  and 
other  parts  of  plants,  and  impart  to  these  their  characteristic 
perfumes  or  flavors.  So  amyl  acetate,  or  banana  oil,  smells  like 
bananas;  methyl  butyrate  C3H7COOCH3  is  known  as  pineapple 
oil;  methyl  salicylate  C6H4(OH)COOCH3  is  the  oil  of  winter- 
green  {G-aultheria procumbens').  Again,  many  of  our  well-known 
fats  and  oils  are  esters.  So  olive  oil  and  cotton-seed  oil-are  essen- 
tially trioleine,  an  ethereal  salt  in  which  the  oleic  acid  radical  is 
united  with  the  glycerine  radical  which  acts  as  base,  thus  : 
(C17H33COO)3*C3H5.  Similarly  in  beef  tallow,  tripalmitine 
(C15H31COO)3-C3H5  and  tristearine  (C17H35COO)3  •  C3H5  are 
found  together  with  some  trioleine.  In  fact,  these  three  esters  in 
which  glycerine  is  the  base  make  up  the  fats.  Trioleine  is  a  liquid 
at  ordinary  temperatures,  whereas  tristearine  and  tripalmitine 
are  solids.  In  mutton  tallow,  which  is  hard,  tristearine  predomi- 
nates, whereas  in  beef  tallow  there  is  more  tripalmitine  and  tri- 
oleine. In  lard,  trioleine  is  still  more  abundant,  which  is  indicated 
by  the  soft,  pasty  consistency  of  the  material. 


HYDROCARBONS   AND   THEIR   DERIVATIVES  253 

On  treating  esters  with  caustic  alkalies,  decomposition  oc- 
curs, thus  :  — 

CH3COOC2H5  +  NaOH  =  CH3COONa  +  C2H5OH. 

A  similar  action  occurs  when  esters  are  boiled  simply  with 
water  :  — 

CH3COOC2H5  +  H20  ^±  CH3COOH  +  C2H5OH. 

In  this  case  we  have  a  typical  instance  of  hydrolysis.  The 
action  is  incomplete,  being  reversible. 

When  fats  are  boiled  with  caustic  soda,  glycerine  and  the 
sodium  salts  of  stearic,  palmitic,  and  oleic  acids  result.  The 
latter  are  soaps.  The  reactions  may  be  expressed  thus  :  — 

(C16H31COO)3C8H6+  3  NaOH  =  3  C16H31COONa+  C8H6(OH)3. 
(C17H36COO)3C3H6  +  3  NaOH  =  3  C1THMCOONa  +  C8H6(OH)3. 
(C17H33COO)3C3H6+3NaOH  =  3C17H33COONa+C8H8(OH)8. 

In  each  case  the  sodium  salts  are  solids  soluble  in  water.  By 
addition  of  common  salt  brine  these  soaps  are  precipitated  or 
"salted  out,"  while  the  glycerine  remains  dissolved.  The  potas- 
sium salts  of  mixtures  of  these  higher  fatty  acids  are  soft  soap, 
and  the  sodium  salts  are  hard  soap.  The  real  nature  of  fats 
became  known  through  the  investigations  of  the  French  chem- 
ist Chevreul. 

The  process  of  decomposing  any  ester  by  means  of  an  alkali  is 
called  saponification.  This  process  is  essentially  the  same  in 
nature,  no  matter  which  ester  is  decomposed. 

When  a  soap  solution  is  mixed  with  hard  water,  a  curdy 
white  precipitate  is  formed  ;  this  is  the  calcium  soap  or  calcium 
salt  of  the  fatty  acid.  Writing  the  reaction  for  sodium  oleate, 
Castile  soap,  and  calcium  sulphate  solution,  we  have  :  — 

2  C17H33COONa  +  CaSO4  =  NaaSO4  +  (C17H33COO)2Ca. 


The  latter  compound  is  the  insoluble  calcium  soap. 

The  cleansing  power  of  soap  depends  upon  the  fact  that  soap 
is  soluble  in  water  due  to  its  content  of  sodium  or  potassium, 
and  that  the  soap  also  coalesces  with  fats  because  of  the  large 
fatty  radical  it  contains.  Thus  soap  loosens  the  greasy  mate- 
rial from  the  skin  or  fabrics,  forming  an  emulsion  which  may 
be  washed  away. 


254  OUTLINES  OF   CHEMISTRY 

Butter  fat  consists  of  about  92  per  cent  of  a  mixture  of  tri- 
oleine,  tripalmitine,  and  tristearine,  and  about  7.7  per  cent  of 
tributyrine,  together  with  smaller  amounts  of  other  glycerides 
that  give  butter  its  characteristic  flavor.  Butter  usually  con- 
tains from  12  to  15  per  cent  water,  also  minor  amounts  of  salt, 
casein,  and  milk  sugar.  The  butter  fat  in  butter  amounts  to 
from  82  to  85  per  cent. 

Nitroglycerine  C3H5(NO)3  is  an  ester  made  by  adding  glyc- 
erine to  a  well-cooled  mixture  of  nitric  and  sulphuric  acid. 
The  latter  serves  as  a  dehydrating  agent.  Nitroglycerine  is 
a  colorless,  odorless,  viscous  liquid  which  explodes  violently 
when  heated  rapidly  to  200°,  or  when  jarred  mechanically. 
Mixed  with  infusorial  earth,  nitroglycerine  forms  a  mass  called 
dynamite,  which  may  be  transported  with  far  less  danger.  On 
treating  nitroglycerine  with  caustic  soda,  sodium  nitrate  and 
glycerine  are  formed. 

Ethers.  — Just  as  the  alcohols  are  the  hydroxides,  so  the  ethers 
are  the  oxides  of  alkyl  radicals.  The  relations  between  water, 
alcohol,  and  ether  on  the  one  hand  and  metallic  hydroxides  and 
oxides  on  the  other  hand  are  illustrated  by  the  following 
formulae :  — 

H2O  water,  NaOH  sodium  hydroxide,  Na2O  sodium  oxide, 
(C2H5)OH  alcohol,  (C2H5)2O  ether. 

Ethers  are  readily  formed  by  treating  alcohols  with  a 
dehydrating  agent  like  sulphuric  acid.  Ordinary  ether  is 
ethyl  oxide.  It  is  made  by  carefully  running  alcohol  into 
concentrated  sulphuric  acid  at  140-145°.  The  reaction  is 
2  C2H5OH  =  (C2H5)2O  +  H2O,  the  water  being  taken  up  by 
the  sulphuric  acid.  For  the  reason  that  sulphuric  acid  is  used 
in  the  manufacture  of  ether,  the  latter  is  often  termed  sulphuric 
ether.  It,  however,  does  not  contain  sulphuric  acid,  for  the 
latter  remains  behind  as  the  ether  distills  off. 

Ethyl  ether  is  a  mobile  liquid  of  agreeable  odor.  It  boils  at 
35°  and  its  specific  gravity  at  0°  is  0.736.  It  is  very  inflamma- 
ble. It  is  used  as  an  anaesthetic  and  also  as  a  solvent  for  fats, 
oils,  and  kindred  substances.  By  using  other  alcohols,  ethers 
of  various  composition  may  be  prepared. 

Sulphur  ethers  are  compounds  in  which  the  oxygen  of  ordi- 
nary ethers  is  replaced  by  sulphur.  Sulphur  ethers  are  vola- 
tile, inflammable  liquids  with  an  extremely  nauseating  odor. 


HYDROCARBONS  AND  THEIR  DERIVATIVES  255 

They  may  be  regarded  as  derived  from  H2S  by  replacing  the 
hydrogen  atoms  by  alkyl  groups;  thus,  we  have  methyl  sul- 
phide (CH3)2S,  ethyl  sulphide  (C2H5)2S,  etc. 

Ketones.  —  Acetone  is  the  simplest  representative  of  a  class 
of  organic  compounds  called  ketones.  It  may  be  prepared  by 
the  dry  distillation  of  sodium  acetate : 

2  CH3COONa  =  Na2CO3  +  (CH3)2CO. 

acetone 

By  heating  the  sodium  salts  of  other  acids,  other  ketones  may 
similarly  be  formed.  The  general  formula  of  a  ketone  is 


Acetone  is  a  colorless,  mobile  liquid  of  agreeable  odor.  Its 
specific  gravity  is  0.792  at  20°.  It  boils  at  56.5°  and  is  rnisci- 
ble  in  all  proportions  with  water  or  alcohol.  Both  ketones 
and  aldehydes  contain  the  carbonyl  group  C  =  O.  In  ketones 
two  radicals  are  combined  with  the  CO  group,  whereas  in  alde- 
hydes one  hydrogen  atom  and  one  radical  are  combined  with  this 
group.  Acetone  is  important  as  a  solvent.  It  commonly  occurs 
in  wood  alcohol,  being  one  of  the  products  of  the  destructive 
distillation  of  wood.  Acetone  is  also  used  in  making  iodoform. 

Carbohydrates. — The  carbohydrates  are  compounds  that  con- 
sist of  carbon,  hydrogen,  and  oxygen,  the  last  two  elements 
being  present  in  the  same  proportions  as  in  water,  whence 
the  name  carbohydrates.  This  is  one  of  the  most  important 
groups  of  compounds,  for  it  contains  the  substances  that  are  found 
in  greatest  abundance  in  the  vegetable  world,  namely,  (1)  the 
sugars,  (2)  the  starches,  (3)  the  celluloses,  (4)  dextrine  and 
gums. 

The  most  important  sugars  may  be  divided  into  two  groups, 
the  monoses,  having  the  empirical  formula  C6H12O6,  and  the 
bioses,  having  the  formula  Ci2H22On.  The  monoses  are  not 
altered  by  dilute  acids,  whereas  the  bioses  are  converted  into 
monoses  by  the  action  of  dilute  acids.  Among  the  important 
monoses  are  (1)  glucose  C6H12O6  or  CH2  •  OH  •  (CH  -  OH)4 •  CHO, 
also  known  as  grape  sugar  or  dextrose;  and  (2)  Icevulose  C6H12O6 
or  CH2(OH).(CHOH)3CO-CH2.OH,  also  called  fructose  or 
fruit  sugar.  From  the  formulae,  it  is  evident  that  these  sugars 
are  polyalcohols  and  that  in  addition  glucose  is  an  aldehyde 
and  Isevulose  a  ketone. 


256  OUTLINES   OF  CHEMISTRY 

Glucose  or  dextrose  occurs  in  the  juice  of  grapes  (whence  its 
name,  grape  sugar)  and  in  many  other  sweet  fruits.  It  is  also 
found  in  honey,  in  the  blood,  and  in  the  urine  of  diabetic 
patients.  Solutions  of  dextrose  turn  the  plane  of  polarized 
light  to  the  right,  whence  the  name  dextrose.  A  Isevo  variety 
has  been  made  in  the  laboratory.  Dextrose  may  be  prepared 
from  cane  sugar  by  the  action  of  dilute  acids.  The  process, 
which  is  called  inversion  of  cane  sugar,  may  be  represented 
thus  :  — 

C12H22On  +  H20  =  C6H1206  +  C6H1206. 

cane  sugar  dextrose  Itevulose 

On  evaporation,  Isevulose  remains  in  solution,  while  dextrose 
separates  out.  Dextrose  crystals  obtained  from  aqueous  solu- 
tions have  the  composition  C6H12O6  +  H2O.  They  melt  at  86°, 
whereas  the  anhydrous  substance  melts  at  146°.  The  action  of 
yeast  readily  converts  dextrose  to  alcohol  and  carbon  dioxide, 
thus  :  — 


Solutions  of  glucose  readily  reduce  hot,  alkaline  solutions  of  cop- 
per sulphate  and  Rochelle  salt,  known  as  Fehling's  solution,  caus- 
ing a  red  precipitate  of  cuprous  oxide  Cu2O  to  form.  This 
test  is  much  used  in  practice,  particularly  in  testing  the  urine  of 
diabetics. 

Glucose  is  only  about  three  fifths  as  sweet  as  cane  sugar. 
Glucose  is  manufactured  in  large  quantities  from  starch  by 
boiling  the  latter  with  dilute  sulphuric  acid,  the  reaction 
being  — 

C6H1005  +  H20  =  C6H1206. 

starch  glucose 

The  acid  is  finally  removed  by  treating  with  calcium  carbonate 
and  filtering  off  the  calcium  sulphate  formed.  In  the  United 
States  large  quantities  of  glucose  are  thus  made  annually  from 
starch  obtained  from  corn. 

Laevulose,  or  fructose,  commonly  occurs  with  glucose  in  fruits 
and  in  honey.  As  already  stated,  it  is  formed  together  with 
glucose,  by  inverting  cane  sugar.  Crystals  of  Isevulose  may  be 
obtained  from  alcoholic  solutions.  They  melt  at  95°.  Solu- 
tions of  Isevulose  turn  the  plane  of  polarized  light  to  the  left, 
whence  the  name  laevulose.  Pure  Isevulose  may  be  obtained 


HYDROCARBONS   AND   THEIR  DERIVATIVES  257 

from  inulin  (a  starch  that  occurs  in  dahlia  roots  and  many  other 
plants)  by  boiling  with  dilute  sulphuric  acid,  thus  :  — 

CH0+H0  =  CH0. 


I052 

inuline  laevulose 


Fructose  also  reduces  Fehling's  solution,  and  by  means  of  yeast 
it  may  be  converted  into  alcohol  and  carbon  dioxide. 

Among  the  most  important  bioses  are  (1)  sucrose,  or  cane 
sugar,  (2)  maltose,  or  malt  sugar,  (3)  lactose,  or  milk  sugar 
Each  of  these  has  the  empirical  formula  C12H22On. 

Cane  sugar,  also  called  saccharose  or  sucrose,  C^H^On,  ^s  very 
widely  distributed  in  the  vegetable  kingdom.  Sugar  cane 
contains  from  15  to  20  per  cent  sucrose,  and  sugar  beets  in 
some  cases  contain  up  to  20  per  cent.  It  is  further  found  in 
the  juice  of  the  sugar  maple,  in  sorghum,  in  the  flowers  of 
plants,  and  in  many  varieties  of  nuts.  Cane  sugar  crystallizes 
in  beautiful  monoclinic  prisms  (rock  candy).  It  is  very  solu- 
ble in  water  ;  1  part  dissolves  in  one  third  its  weight  of  water 
at  room  temperatures.  At  160°  sucrose  melts,  and  on  cooling 
it  solidifies  to  an  amorphous  glassy  mass  called  barley  sugar, 
which  on  long  standing  again  becomes  crystalline.  On  heat- 
ing cane  sugar  somewhat  over  200°,  water  is  given  off  and  a 
brown  substance  called  caramel  is  formed.  As  already  stated 
under  dextrose,  sucrose  yields  dextrose  and  Isevulose  on  hy- 
drolysis, the  biose  being  thus  split  into  two  monoses.  A  cane 
sugar  solution  does  not  reduce  Fehling's  solution.  It  does  not 
readily  ferment  with  yeast.  The  latter,  however,  contains  an 
enzyme  called  invertase,  which  inverts  cane  sugar,  and  the 
fructose  and  glucose  thus  formed  are  then  converted  into 
alcohol  and  carbon  dioxide  by  fermentation. 

Nearly  ten  million  tons  of  sugar  are  produced  annually  from 
sugar  cane  and  sugar  beets.  The  process  of  preparing  sugar 
commercially  consists  of  expressing  the  juice  from  the  cane  or 
beets  and  boiling  this  with  about  a  1  per  cent  solution  of  milk 
of  lime  to  neutralize  acids,  coagulate  vegetable  albuminous 
matter,  and  prevent  fermentation.  The  solution  so  obtained 
is  then  treated  with  carbon  dioxide  to  remove  the  excess  of 
lime,  and  filtered  through  bone  black  to  decolorize  it.  It  is 
then  concentrated  by  evaporation  in  a  partial  vacuum  in  so- 
called  vacuum  pans,  which  are  heated  with  steam.  The  crys- 
tals which  are  deposited  on  cooling  are  freed  from  the  adhering 


258  .  OUTLINES   OF   CHEMISTRY 

brown  mother  liquor  by  centrifugal  force  in  the  "  centrifugals." 
The  latter  consist  of  sieves  rapidly  rotated  by  machinery.  The 
mother  liquor  is  hurled  through  the  meshes  of  these  sieves,  and 
the  crystals  remain  behind  almost  dry.  The  drying  is  com- 
pleted with  the  aid  of  heat.  The  mother  liquor,  from  which 
crystals  will  no  longer  form  because  of  the  presence  of  various 
impurities,  is  called  molasses.  The  residue  of  the  cane  or  beets 
from  which  the  juice  has  been  extracted  is  called  the  begasse. 
It  is  made  into  paper  or  used  as  fuel. 

With  lime,  cane  sugar  forms  calcium  sucrate  C12H22On  •  CaO, 
which  is  soluble  in  water.  The  analogous  strontium  sucrate 
C12H22On  •  SrO  is  used  as  a  means  of  recovering  further  crystal- 
lizable  sugar  from  molasses. 

Cane  sugar  solutions  rotate  the  plane  of  polarized  light  to 
the  right.  The  specific  rotatory  power  of  cane  sugar  for  sodium 
light  is  4-  66.5°,  while  that  of  dextrose  is  +  52.7°,  and  of  Isevu- 
lose  —  93°.  Consequently,  solutions  of  invert  sugar,  which 
contain  equal  parts  of  dextrose  and  laevulose,  always  exhibit 
laevorotatory  power. 

Maltose  C12H22On  4-  H2O  crystallizes  with  one  molecule  of 
water,  as  its  formula  indicates.  It  may  be  formed  by  the  action 
of  diastase,  a  ferment  contained  in  malt,  upon  starch.  Maltose 
is  of  importance  as  an  intermediate  product  in  the  manufacture 
of  alcohol  from  starch.  Maltose  is  strongly  dextrorotatory,  its 
specific «rotatory  power  being  -f- 137°.  On  boiling  it  with  dilute 
mineral  acids,  dextrose  only  is  produced,  showing  that  maltose 
is  an  anhydride  of  dextrose.  Unlike  cane  sugar,  maltose  reduces 
Fehling's  solution,  and  readily  ferments  with  yeast. 

Lactose  C12H22()n  4-  H2O,  milk  sugar,  is  found  in  milk.  Like 
maltose,  its  crystals  contain  one  molecule  of  water.  It  is  not 
as  sweet  as  cane  sugar,  and  is  much  less  soluble  in  water;  1 
part  dissolves  in  6  parts  of  water.  Lactose  is  dextrorotatory, 
its  specific  rotation  being  4-  52.5°.  On  boiling  with  dilute  min- 
eral acids,  it  yields  glucose  and  another  monose  called  galactose 
C6H12O6,  thus :  — 

Ci2H22On  4-  H20  =  C6H1206  +  C6H1206. 

lactose  glucose  galactose 

Lactose  reduces  Fehling's  solution.  But  the  reduction  is  less  rapid 
than  when  glucose  is  used.  Pure  yeast  does  not  produce  fermen- 
tation of  milk  sugar ;  but  ordinary  yeast  acts  upon  milk  sugar 


HYDROCARBONS   AND   THEIR  DERIVATIVES  259 

as  it  does  upon  cane  sugar.  The  products  formed  are  alcohol 
and  lactic  acid.  Cow's  milk  contains  nearly  5  per  cent  milk 
sugar. 

Fermentation  and  Enzymes.  —  Under  fermentation  are  classed 
a  large  number  of  chemical  processes  that  are  produced  directly 
or  indirectly  by  organisms  commonly  termed  ferments.  These 
organisms  are  yeasts,  molds,  or  fungi,  and  bacteria.  They  se- 
crete complex  compounds  called  enzymes,  the  chemical  nature 
of  which  is  not  yet  understood.  They  seem  to  be  closely  related 
to  the  albumins  and  peptones  (which  see).  In  the  presence  of 
these  enzymes,  which  are  also  called  unorganized  ferments,  fer- 
mentation takes  place,  each  enzyme  causing  its  own  particular 
chemical  change,  which  commonly  progresses  at  room  tempera- 
tures, being  accomplished  with  evolution  of  heat  and  frequently 
with  liberation  of  gas.  As  a  rule  the  action  is  checked  by  either 
raising  or  lowering  the  temperature  materially,  also  by  the 
addition  of  various  poisons.  The  action  of  the  enzymes  is  often 
termed  catalytic,  for  a  small  amount  of  enzyme  may  produce  a 
large  amount  of  change  without  being  itself  seemingly  affected. 
Among  the  common  enzymes  are :  zymase,  contained  in  yeast 
cells,  producing  alcoholic  fermentation ;  malt  diastase,  contained 
in  malt,  converting  starch  into  malt  sugar ;  invertase,  contained 
in  ordinary  yeast,  inverting  cane  sugar  and  lactose ;  emulsin, 
contained  in  bitter  almonds ;  pepsin,  contained  in  the  stomach 
juices,  aiding  in  the  digestion  of  albumins ;  trypsin,  contained 
in  the  intestinal  juices,  aiding  in  so-called  tryptic  digestion. 

Starch  and  Dextrine.  —  Starch  (CgH^Og)^  is  a  carbohydrate 
found  in  large  quantities  in  all  plants,  particularly  in  tubers 
like  potatoes,  and  in  grains  like  corn,  rye,  oats,  wheat,  rice,  etc. 
Starch  is  insoluble  in  water  and  occurs  as  a  nodular  deposit  of 
varying  sizes  and  forms  in  different  plant  cells.  Figure  95  shows 
grains  of  potato  starch  as  they  appear  under  the  microscope. 
Figure  96  shows  grains  of  wheat  starch,  and  Figure  97  shows 
grains  of  corn  starch.  Figure  98  shows  potato  starch  grains  as 
they  look  in  polarized  light.  Starch  is  prepared  mainly  from 
potatoes  and  corn,  the  former  being  commonly  used  in  Europe 
and  the  latter  in  America.  The  potatoes  or  grains  are  ground 
so  as  to  break  up  the  plant  cells  and  lay  the  starch  granules 
bare.  These  are  then  washed  out  with  water,  forming  a  thin 
milky  paste  which  is  passed  through  fine  sieves  that  retain  the 


260 


OUTLINES  OF  CHEMISTRY 


pulp.  The  starch  is  allowed  to  settle,  the  water  is  drained 
off,  and  the  remaining  material  is  dried  at  low  temperatures. 
Though  starch  is  insoluble  in  cold  water,  its  granules  swell  up 
and  burst  when  treated  with  boiling  water,  and  the  contents  of 


FIG.  95. 


FIG.  96. 


the  cells,  the  granulose,  dissolves,  whereas  the  cell  walls  remain 
undissolved  and  can  be  filtered  off.  From  the  filtrate,  the  gran- 
ulose or  soluble  starch  may  be  precipitated  by  adding  alcohol. 
Starch  paste  is  prepared  by  grinding  starch  with  a  little  cold 


FIG.  97. 


FIG.  98. 


water  and  then  adding  boiling  water  with  constant  stirring. 
Thus  a  semi-transparent,  gelatinous  mass  results,  which  is  used 
as  an  adhesive.  It  is  also  employed  in  the  laundry  for  stiffen- 
ing clothes,  the  hot  sad  iron  converting  the  starch  into  dextrine, 


HYDROCARBONS  AND   THEIR  DERIVATIVES  261 

which  covers  the  linen,  imparting  to  it  a  characteristic  luster. 
Starch  has  the  empirical  formula  C6H10O5.  Its  molecular 
weight  is  unknown,  but  it  is  probably  some  rather  high  multi- 
ple of  the  formula  given.  Starch  is  an  exceedingly  important 
article  of  food.  Wheat  flour,  for  instance,  consists  of  about  70 
per  cent  starch,  10  per  cent  gluten,  a  sticky  nitrogenous  sub- 
stance closely  allied  to  albumins  like  white  of  egg,  and  minor 
amounts  of  water,  sugar,  and  inorganic  material.  Gluten,  like 
starch,  is  valuable  as  food  material. 

On  heating  starch  to  about  210°  it  is  converted  into  dextrine 
(C6H10O5),  which  is  also  obtained  as  an  intermediate  product 
in  the  conversion  of  starch  to  glucose  by  means  of  dilute  sul- 
phuric acid.  Dextrine  is  a  colorless,  amorphous  substance  that 
has  a  strong  adhesive  property.  It  is  consequently  frequently 
used  as  a  cheap  gum. 

Cellulose.  —  This  carbohydrate  has  the  same  empirical  formula 
as  starch  (Cgtl^Og)^.  Its  molecular  weight  is  unknown,  but 
it  is  probably  rather  high.  Cellulose  is  very  widely  distributed 
in  nature,  constituting  the  material  out  of  which  the  cell  walls 
of  plants  are  made.  Thus  wood,  cotton,  linen,  hemp,  flax,  etc., 
when  freed  from  the  mineral  matter  they  contain,  are  almost  pure 
cellulose.  Filter  paper  after  extraction  with  hydrochloric  and 
hydrofluoric  acids  and  washing  with  water  is  a  nearly  pure  form 
of  cellulose.  Cellulose  is  insoluble  in  water  but  soluble  in  a 
solution  of  copper  in  strong  ammonia  water,  which  is  known  as 
Schweitzer's  reagent.  From  this  solution  cellulose  may  be  pre- 
cipitated by  means  of  acids  in  the  form  of  a  gelatinous  mass  that 
may  be  filtered  off  and  dried  to  an  amorphous  powder.  Cellur 
lose  gradually  dissolves  in  concentrated  sulphuric  acid.  On 
diluting  the  solution  and  boiling,  dextrine  is  formed,  which  is 
finally  converted  into  glucose.  Thus  wood  may  be  changed  to 
glucose,  from  which  in  turn  alcohol  may  be  obtained  by  fermen- 
tation. Paper  consists  essentially  of  thin  sheets  of  fibers  of 
cellulose  matted  together. 

On  heating  cotton  with  a  mixture  of  nitric  and  sulphuric 
acids,  nitrates  of  cellulose  are  formed.  This  action  is  similar 
to  the  formation  of  glycerine  nitrate,  i.e.  nitroglycerine. 
The  composition  of  cellulose  nitrates  varies  according  to  the 
length  of  time  that  the  nitric  acid  has  acted  upon  the  cellu- 
lose, and  the  concentration  of  the  acid.  Cellulose  hexanitrate 


262  OUTLINES   OF   CHEMISTRY 


6O4  *s  calle(l  nitrocellulose,  or  gun  cotton.  It  does 
not  dissolve  in  a  mixture  of  alcohol  and  ether.  Gun  cotton 
looks  like  ordinary  cotton,  but  it  is  not  as  soft  to  the  touch.  It 
burns  rapidly  and  quietly,  producing  no  smoke.  By  means  of 
fulminating  mercury  it  can  be  made  to  explode  violently.  It 
is  used  as  an  explosive,  frequently  together  with  nitroglycerine. 
Gun  cotton  and  nitroglycerine  are  made  into  threads  with  the 
aid  of  acetone  and  vaseline  and  are  thus  used  as  smokeless  gun- 
powder. Celluloid  consists  of  camphor  and  gun  cotton.  It  is 
not  explosive,  but  it  burns  readily.  The  tetra-  and  petranitrates 
of  cellulose  C12H16(NO3)4O6  and  C12H15(NO3)5O5  readily  dis- 
solve in  a  mixture  of  alcohol  and  ether.  This  solution  is  known 
as  collodion.  It  is  used  in  photography  for  making  films,  and  in 
surgery  for  protecting  wounds  from  the  air,  for  when  collodion 
solution  is  poured  out  in  thin  layers,  the  alcohol  and  ether 
evaporate,  leaving  a  tough,  amorphous,  transparent  layer  of 
nitrocellulose.  On  treating  nitrocellulose  with  caustic  soda, 
cellulose  and  sodium  nitrate  are  formed,  showing  that  nitrocettu- 
lose  is  a  nitrate  of  cellulose. 

Nitrobenzene  C6H5NO2  is  a  light  yellow  oil  formed  by  the 
action  of  nitric  acid  on  benzene.  Its  boiling  point  is  208°.  It 
is  used  as  a  perfume  for  laundry  soaps,  under  the  name  oil  of 
mirbane.  By  reducing  nitrobenzene  with  nascent  hydrogen, 
aniline  C6H5NH2  is  formed.  Aniline  is  a  base  which  forms 
salts  like  ammonia.  Thus,  C6H5-NH2-HC1  is  aniline  hydro- 
chloride.  Aniline  is  a  nearly  colorless  liquid  which  boils  at 
189°.  On  exposure  to  the  air,  it  soon  turns  brown.  Aniline 
may  be  regarded  as  ammonia  in  which  one  hydrogen  atom  has 
been  replaced  by  the  phenyl  group,  C6H5.  Many  similar  sub- 
stituted ammonias  are  known;  they  are  called  amines.  So  we 
have  methyl  amine  CH3-NH2,  ethyl  amine  C2H5  •  NH2,  propyl 
amine  C3H7-NH2,  etc.  In  general,  their  chemical  behavior  is 
similar  to  that  of  ammonia.  The  aniline  dyes  are  deriva- 
tives of  pararosaniline  HO  •  C(C6H4NH2)3  and  rosaniline 
HO  •  C(C6H4NH2)2  .  C6H3  •  CH3  .  NH2.  The  hydrochloride  of 
the  latter  is  the  dyestuff  magenta.  Many  aniline  dyes  are 
known.  They  represent  very  many  beautiful  colors.  The  ani- 
line dyestuff  industry  is  an  important  one.  It  will  be  noted 
that  the  dyes  are  all  derived  from  benzene,  which  is  obtained 
from  coal  tar,  whence  the  name  coal-tar  dye. 


HYDROCARBONS   AND   THEIR   DERIVATIVES  263 

Alkaloids  are  complex  basic  substances  that  occur  in  plants. 
These  compounds  are  so  named  because  they  form  salts  with 
acids,  thus  playing  the  role  of  alkalies.  The  base  pyridine 
C6H5N  is  a  colorless,  odoriferous  liquid  boiling  at  115°,  which 
is  found  in  coal  tar  and  in  the  products  of  the  dry  distilla- 
tion of  green  bones.  Quinoline  C9H7N  occurs  with  pyridine. 
Many  of  the  plant  alkaloids  are  related  to  these  two  bases. 
Nicotine  C10H14N2  is  found  in  tobacco.  It  is  a  poisonous  liquid, 
boiling  at  247°.  Atropine  C17H23O3N  occurs  in  nightshade, 
Atropa  belladonna.  It  forms  very  poisonous  crystals  which  melt 
at  115.5°.  It  is  used  by  oculists  to  cause  expansion  of  the 
pupil  of  the  eye.  Cocaine  C17H21O4N  is  found  in  coca  leaves. 
It  consists  of  crystals  that  melt  at  98°.  It  is  used  as  a  local 
anaesthetic.  Quinine  C20H24O2N2  and  cinchonine  C19H22O2N2  are 
found  in  Peruvian  bark,  which  also  contains  other  alkaloids. 
Quinine  is  usually  administered  in  form  of  the  sulphate.  It  is 
a  specific  for  malarial  fever.  Morphine  C17H19O3N,  narco- 
tine  C22H23O7N,  and  codeine  C18H21NO3  occur  in  opium,  which 
is  the  dried  sap  of  the  unripe  seed  pods  of  the  opium  poppy. 
Morphine  is  a  crystalline  powder  which  was  first  isolated  in 
1806.  It  induces  sleep.  As  a  rule  it  is  administered  in  form 
of  the  hydrochloride.  Strychnine  C21H22O2N2  and  brucine 
^23^26^2^4  occur  in  nux  vomica.  These  alkaloids  are  very  poison- 
ous, causing  death  accompanied  by  muscular  contraction  and  rigor. 
Strychnine  forms  crystals  that  melt  at  265°.  They  are  nearly 
insoluble  in  water.  Strychnine  compounds  have  a  very  bitter 
taste.  In  minute  quantities  they  are  often  prescribed  in  medi- 
cine as  a  tonic. 

Proteins.  —  These  are  very  important  compounds,  consisting  of 
carbon,  hydrogen,  nitrogen,  oxygen,  and  sulphur,  which  make  up  a 
large  share  of  the  bodies  of  animals  and  also  play  an  important 
role  in  plants.  Exclusive  of  water,  fats,  and  mineral  constitu- 
ents, the  animal  matter  consists  of  proteins,  formerly  called 
proteids.  Without  proteins  as  food,  animals  will  finally  die. 
The  chemical  structure  of  the  proteins  is  very  complicated. 
On  the  average,  proteins  have  about  the  following  com- 
position :  — 


264  OUTLINES  OF  CHEMISTRY 

Carbon  50.5  to  54.5  per  cent 

Hydrogen  6.5  to    7.3  per  cent 

Nitrogen  15.0  to  17.7  per  cent 

Oxygen  21.0  to  24.0  per  cent 

Sulphur  0.3  to    2.3  per  cent 

Phosphorus  0.4  to    0.8  per  cent 

Nucleoproteins  may  contain  from  5  to  6  per  cent  phosphorus. 

Among  some  of  the  common  protein  substances  may  be 
mentioned  the  albumins,  as  they  occur  for  instance  in  serum, 
in  eggs,  in  milk,  in  muscular  tissues,  and  in  leguminous  and 
other  seeds.  Albumins  are  coagulated  by  heat  and  by  acids. 
Albuminoids  are  closely  related  to  albuminous  bodies.  Among 
the  albuminoids  are  gelatine,  elastine,  and  keratine.  Elastine 
enters  into  the  composition  of  connective  tissues,  while  kera- 
tine is  the  main  substance  found  in  hoofs,  horns,  feathers,  hair, 
nails,  and  the  epidermis.  By  the  action  of  pepsin  albumins 
are  converted  into  peptones,  which  process  involves  an  addition 
of  the  elements  of  water  to  the  molecule. 

The  molecular  weight  of  the  protein  molecule  as  determined 
by  the  freezing  point  method  (which  see)  is  approximately 
15,000.  Proteins  are  consequently  regarded  as  colloidal  sub- 
stances. The  fact  that  they  diffuse  very  slowly  in  solutions 
also  accords  with  this  view.  The  idea  that  proteins  are  very 
complicated  bodies  comes  from  the  fact  that  a  large  number  of 
products  may  be  obtained  from  them  by  treatment  with  various 
reagents  like  acids,  alkalies,  oxidizing  agents,  etc.  The  study 
of  proteins  has  of  recent  3rears  been  prosecuted  with  special 
success  by  Professor  Emil  Fischer  of  the  University  of  Berlin, 
to  whom  we  also  owe  much  of  our  knowledge  of  the  constitu- 
tion of  sugars. 

When  protein  substakces,  like  meat,  fish,  etc.,  putrefy,  pto- 
maines are  often  formed.  These  are  basic  substances  which 
act  like  the  alkaloids  in  many  respects.  They  are  poisonous. 
Among  them  are  putrescine  C4H8(NH2)2  and  cadaverine 
C5H10(NH2)2.  Illness  from  eating  spoiled  meat  is  commonly 
due  to  ptomaine  poisoning. 


CHAPTER   XVI 

ILLUMINATING   GAS   AND   FLAMES 

Illuminating  Gas.  —  Illuminating  gas  is  often  produced  by 
the  destructive  distillation  of  soft  coal.  During  this  process 
there  are  formed  coal  gas,  coal  tar,  and  water  containing 
ammonia  and  other  products  in  solution.  Coal  gas  consists 
mainly  of  hydrogen,  carbon  monoxide,  and  methane,  together 
with  a  small  amount  of  higher  hydrocarbons  known  as  illumi- 
nants.  Impurities  like  nitrogen,  carbon  dioxide,  and  hydrogen 
sulphide  are  also  present  in  small  quantities.  The  tar  from 
the  gas  works  is  distilled,  and  the  benzene  obtained  therefrom 
is  used  in  manufacturing  dyestuffs  and  many  medicinal  prep- 
arations. The  remaining  tar  is  then  employed  for  roofing, 
making  artificial  asphalt,  black  varnishes,  etc.  The  coke  from 
the  gas  retorts  serves  as  fuel.  The  ammoniacal  gas  liquors 
serve  as  a  source  of  ammonium  salts.  These  liquors  are  neu- 
tralized with  sulphuric  acid,  and  the  ammonium  sulphate  so 
obtained  is  used  as  a  fertilizer  or  converted  into  ammonia 
water  by  treatment  with  lime  and  absorption  of  the  gas  lib- 
erated. The  coal  gas  is  treated  with  lime  to  remove  carbon 
dioxide  and  sulphides,  particularly  hydrogen  sulphide,  which 
is  always  present. 

In  the  process  of  manufacture,  coal  gas  passes  from  the 
retort  R  (Fig.  99),  through  the  condensers  (7,  in  which  the  tar 
and  ammoniacal  liquors  are  condensed,  into  the  scrubber  S, 
where  it  is  washed  with  water.  The  gas  then  goes  through 
layers  of  lime  or  oxides  of  iron,  or  both,  which  are  contained 
in  the  purifiers  P.  This  removes  the  hydrogen  sulphide. 
Finally  the  gas  enters  the  holder  H,  from  which  it  is  distrib- 
uted through  the  mains. 

Not  all  illuminating  gas  is  coal  gas.  Much  illuminating  gas  is 
made  in  America  from  petroleum.  By  allowing  the  latter  to 
come  into  contact  with  the  walls  of  a  chamber  lined  with  fire 
bricks  heated  to  about  1000°  C.  the  heavy  hydrocarbons  are 

265 


266 


OUTLINES   OF   CHEMISTRY 


ILLUMINATING   GAS   AND   FLAMES  267 

decomposed  into  simple  hydrocarbons  that  contain  fewer  car- 
bon atoms  in  the  molecule,  and  are  consequently  gaseous  at 
ordinary  temperatures.  This  process  is  called  "cracking"  the 
petroleum  oils.  Pintsch  gas  is  made  entirely  by  this  process. 
It  has  very  high  illuminating  power  and  is  used  for  lighting 
railway  cars,  for  which  purpose  it  is  compressed,  with  moderate 
pressure,  in  steel  cylinders.  A  special  burner  must  be  used 
with  this  gas  to  prevent  smoking  of  the  flame.  Oil  gas  thus 
made  is  rich  in  benzene  and  the  hydrocarbons  of  the  ethylene  and 
acetylene  series.  These  give  it  its  high  illuminating  power, 
and  are  consequently  called  illuminants.  It  is  now  common 
practice  to  increase  the  illuminating  power  of  coal  gas  by 
"  enriching  "  the  gas  by  means  of  petroleum  oils  added  directly 
to  the  coal  in  the  retorts,  or  by  adding  oil  gas  to  the  coal  gas. 
It  will  be  recalled  that  water  gas  consists  of  hydrogen  and 
carbon  monoxide,  and  consequently  burns  with  a  flame  that  is 
nearly  non-luminous.  To  make  it  serve  for  illuminating  pur- 
poses, it  is  enriched,  or  "  carbureted,"  by  means  of  oil  gas. 
To  accomplish  this  the  water  gas  is  passed  through  heated 
chambers  into  which  petroleum  oil  is  run  and  cracked  as  above 
stated,  thus  furnishing  the  necessary  illuminants.  Sometimes 
the  oil  gas  is  made  separately  and  then  added  to  the  water  gas 
in  proper  proportion.  Enriched  water  gas  is  much  used  in  the 
large  cities  of  the  United  States. 

Coal  gas  was  first  used  for  house  illumination  in  1792,  by 
William  Murdock  in  London,  where  it  was  employed  for  street 
lighting  in  181/2.  Three  years  later  it  was  also  used  for  this 
purpose  in  Paris.  However,  the  fact  that  a  combustible  gas  is 
evolved  from  coal  when  it  is  heated,  was  discovered  as  early  as 
the  year  1680  by  Becher. 

The  illuminating  power  of  a  gas  is  usually  expressed  in 
candle  power.  The  gas  is  burned  from  a  burner  consuming 
5  cu.  ft.  per  hour,  and  this  light  is  compared  with  that  of  a 
standard  candle  by  means  of  a  photometer. 

The  composition  of  gases  used  for  illuminating  purposes 
varies  considerably.  A  general  idea  of  the  composition  of 
illuminating  gas  may,  however,  be  obtained  from  the  fol- 
lowing table,  from  the  reports  of  the  Massachusetts  State  Gas 
Inspectors,  which  presents  average  results  in  per  cent  by 
volume  :  — 


268 


OUTLINES   OF   CHEMISTRY 


COAL  GAS 

CARBURETTED 
WATER  GAS 

OIL  GAS 

Candle  power      .... 

17.5 

25. 

65. 

Illuminants     

50 

16  6 

450 

Methane     

345 

19  8 

38  8 

Hydrogen       

490 

32  1 

14  6  l 

Carbon  monoxide    .     .     . 
Nitrogen 

7.2 
32 

26.1 
2  4 

1  l 

Oxvsren 

Carbon  dioxide  .... 

1.1 

3.0 

— 

Flame.  —  Flames  are  produced  by  burning  gases.     Whenever 
a  continuous  stream  of  one  gas  issues  into  an  atmosphere  of 
another  gas  upon  which  it  acts  chemically,  producing  a  suffi- 
cient rise  of  temperature  and  a  certain  degree  of  luminosity, 
we  have  at  the  surface  of  contact  of  the  gases  the  phenomenon 
of  flame.     Solids  like  charcoal  and  coke  burn  practically  with- 
out flame.     In  the  kerosene  lamp  the 
wick  draws  up  the  oil  and  the  latter  is 
converted  into  gas  by  heat,  and  thus 
it  is  really  the  gas  that  barns  in  the 
flame.     Similarly  in   the    flame   of   a 
candle  it  is  gas  that  burns ;  for  the 
wax  or  tallow  is  melted  by  the  heat, 
and  the   liquid  is  drawn  up  by  the 
wick,  being  converted  to  gas  before 
it  is  burned.     This  is  well  illustrated 
by  the  experiment  represented  in  Fig. 
100,  in  which  some  of  the  unconsumed 
gas  from  the  inner  part  of  the  flame 

is  conducted  off  by  means  of  the  glass  tube,  at  the  end  of  which 
it  has  been  lighted. 

Oxygen  is  said  to  support  the  combustion  of  a  gas  that  burns 
in  the  air.  Obviously,  however,  combustion  is  really  the  mutual 
interaction  of  the  two  gases.  So,  for  instance,  coal  gas  can  be 
burned  in  the  form  of  a  flame  in  the  air,  or  air  can  be  burned  in 
the  form  of  a  flame  in  an  atmosphere  of  coal  gas.  This  is  well 
illustrated  by  the  experiment  shown  in  Figs.  101,  102,  and  103. 

1  This  figure  represents  ethane  and  not  hydrogen.  The  latter  is  not 
present  in  oil  gas. 


FIG.  100. 


ILLUMINATING  GAS  AND  FLAMES 


269 


Illuminating  gas  is  passed  into  the  apparatus  (Fig.  101),  through 
the  tube  at  the  left,  and  then  the  gas  is  ignited  at  the  end 
of  the  other  tube  at  the  right,  as  shown.  Now  the  lid  on  top  of 
the  apparatus  is  opened  (Fig.  102),  whereupon  the  draft  created 
causes  the  flame  to  strike  inward  and  burn  inside  of  the  large 
tube,  as  shown  in  the  figure.  This  is  now  a  jet  of  air  burning 
in  an  atmosphere  of  illuminating  gas.  The  gas  issuing  from 
the  opening  created  by  removing  the  lid  may  be  lit  as  shown, 
and  thus  we  have  the  upper  flame  of  illuminating  gas  burning 


FIG.  101. 


FIG.  102. 


FIG.  103. 


in  the  air,  and  at  the  same  time  the  inner  flame  of  air  burning 
in  the  atmosphere  of  the  gas.  This  is  commonly  spoken  of  as 
the  reverse  flame.  That  it  is  actually  air  that  is  burning  in 
the  illuminating  gas  may  be  demonstrated  by  passing  a  small 
gas  flame  up  into  the  inside  of  the  flame  of  air  burning  in  the 
illuminating  gas,  as  shown  in  Fig.  103 ;  the  small  gas  jet  burns 
in  the  inner  cone  of  the  flame. 

It  will  be  recalled  that  a  jet  of  hydrogen  will  burn  in  oxygen 
and  a  jet  of  oxygen  in  an  atmosphere  of  hydrogen ;  and  that  a 
jet  of  chlorine  will  burn  in  hydrogen  and  a  jet  of  hydrogen  in 
an  atmosphere  of  chlorine.  These  experiments  also  demon- 
strate further  the  relative  nature  of  combustion.  The  latter 
experiment  further  shows  that  we  may  have  flames  due  to 
chemical  changes  other  than  oxidation. 


270 


OUTLINES   OF  CHEMISTRY 


FIG.  104. 


The  fact  that  oxygen  will  burn  in  coal  gas  is  readily  demon- 
strated by  the  experiment  illustrated  in  Fig.  104.  Chlorate  of 
potassium  heated  in  a  deflagrating  spoon 
till  oxygen  is  evolved  freely  is  introduced 
into  the  atmosphere  of  coal  gas  as  shown. 
The  oxygen  burns  brilliantly,  the  flame 
being  colored  purple  by  the  presence  of 
potassium.  The  gas  should  be  ignited  at 
the  upper  orifice,  as  shown 
in  the  figure,  before  in- 
troducing the  deflagrat- 
ing spoon. 

Luminosity  of  Flame.  — 
Flames  are  either  lumi- 
nous or  non-luminous. 
Thus,  hydrogen  burns 
with  a  non-luminous 
flame.  This  may,  how- 
ever, be  made  luminous 
by  blowing  fine  solid 
particles,  like  soot,  for 
example,  into  the  flame. 
These  fine  particles  be- 
come heated  to  incandes- 
cence and  emit  light. 
Similarly,  a  platinum  wire 

introduced    into    the    hydrogen    flame    soon 

becomes   hot   and   gives   off   light.       Carbon 

monoxide  burns  with  a  blue  flame,  which  can 

readily  be  made  luminous  by  the  introduction 

of  fine  particles  of  carbon.     Figure  105  shows 

a  convenient  apparatus  for  accomplishing  this. 

Carbon  monoxide  is  passed  through  the  ap- 
paratus, the  cock  at  the  left  being  open,  and 

ignited.     We  now  have  the  characteristic  blue 

flame.     If  the  cock  at  the  right  is  then  opened 

and  the  one  at  the  left  closed,  thus  causing 

the  gas  to  pass  through  the  cotton  saturated 

with  benzene  contained  in  the  right  limb,  the  flame  becomes 

luminous,  for  the  heavy  hydrocarbon  vapors  carried  into  the 


FIG.  105. 


ILLUMINATING   GAS   AND   FLAMES 


271 


flame  lose  their  hydrogen  at  the  high  temperature  of  the  flame, 
thus  setting  particles  of  carbon  free  which  are  heated  to  incan- 
descence and  so  give  the  flame  luminosity.  By  proper  manipu- 
lation of  the  two  cocks,  the  flame  may  be  rendered  luminous  or 
non-luminous  alternately.  That  all  luminous  gas  flames,  in- 
cluding those  of  oil  lamps  and  candles,  contain  carbon  particles 
in  suspension  is  readily  de- 
monstrated by  the  fact  that  a 
layer  of  soot  is  deposited  upon 
glass  or  porcelain  held  in 
such  flames.  According  to 
Lewes,  acetylene  is  formed 
in  all  luminous  flames,  and  it 
would  seem  probable  that  the 
luminosity  is  due  to  the  pres- 
ence of  this  gas,  which  breaks 
up,  yielding  carbon  particles 
that  then  become  heated  to 
incandescence.  At  any  rate, 
it  is  certain  that  the  lumi- 
nosity of  the  flame  of  illumi- 
nating gas  is  due  to  the 
presence  of  benzene  and 
gases  of  the  ethylene  and 
acetylene  series,  which  are 
commonly  called  the  illumi- 
nants.  It  is  well  known  that 
at  high  temperatures  these 
decompose,  yielding  acetylene 
as  one  of  the  products. 

When  a  jet  of  illuminating 
gas  is  mixed  with  air  and 
then  ignited,  a  non-luminous 
flame  is  obtained.  This  is  the  principle  used  in  the  Bunsen 
burner.  Figure  106  shows  a  simple  apparatus  for  illustrating 
this  principle.  Gas  issues  through  the  opening  of  the  small 
tube,  over  which  the  larger  tube  is  placed  as  shown.  The  gas 
passing  upward  creates  a  current  of  air  which  enters  the  larger 
tube,  as  shown  by  the  arrows,  and  mixes  with  the  gas  in  the 
large  tube  at  the  end  of  which  the  mixture  is  lighted,  yielding 


^ 


// 


£. 


\ 


FIG.  106. 


272  OUTLINES  OF  CHEMISTRY 

a  blue,  practically  non-luminous  flame.  It  is  commonly  held 
that  this  non-luminosity  is  due  to  the  oxygen,  which  causes 
complete  combustion  of  the  carbon  particles. 

The  Bunsen  burner  is  used  in  various  forms  in  the  laboratory. 
In  principle,  gas  stoves  and  furnaces  are  all  Bunsen  burners. 
By  supplying  compressed  air  by  means  of  a  bellows  or  other- 
wise, the  blowpipe  or  blast  flame  is  produced.  This  additional 
supply  of  air  insures  more  rapid  and  more  complete  combus- 
tion, and  consequently  a  higher  temperature  is  obtained.  The 
ordinary  form  of  blast  lamp  is  similar  to  the  oxy hydrogen 
lamp,  except  that  coal  gas  and  air  are  used  instead  of  hydrogen 
and  oxygen. 

The  flame  of  an  alcohol  lamp  is  non-luminous  because  alcohol 
contains  some  oxygen,  and  this,  together  with  that  supplied  by 
the  air,  is  sufficient  to  secure  complete  combustion  of  the  par- 
ticles of  carbon.  For  similar  reasons  ether  burns  with  a  prac- 
tically non-luminous  flame.  As  the  carbon  content  of  com- 
pounds increases,  however,  the  flames  with  which  the  compounds 
burn  become  luminous,  and  even  sooty. 

It  will  be  recalled  that  when  calcium  oxide  was  introduced 
into  the  oxyhydrogen  flame,  the  lime  was  heated  to  a  tempera- 
ture at  which  it  emitted  a  brilliant  white  light.  The  brilliant 
light  formed  when  magnesium  or  phosphorus  burn  in  the  air 
or  in  oxygen  is  similarly  caused  by  the  incandescence  of  the 
particles  of  MgO  and  P2O5  that  are  formed  during  the  com- 
bustion. 

In  the  Welsbach  light  this  principle  is  used  by  hanging  over 
the  flame  of  a  Bunsen  burner  a  mantle  consisting  of  a  network 
made  of  99  per  cent  thorium  dioxide  and  1  per  cent  cerium 
dioxide.  Thus  a  strong,  brilliant,  white  light  is  produced  with 
a  relatively  low  consumption  of  gas.  Other  oxides  will  also 
serve,  but  they  do  not  yield  nearly  as  good  results  in  practice. 
The  oxides  mentioned,  when  used  in  the  proportion  indicated, 
have  been  found  to  give  the  strongest  light. 

The  Structure  of  Flame.  —  Taking  as  a  typical  luminous 
flame  that  of  a  candle  (Fig.  107),  we  see  that  it  exhibits  three 
distinct  zones.  The  inner  zone  A  is  dark  and  non-luminous. 
It  consists  of  the  gases  produced  by  the  decomposition  and 
volatilization  of  the  material  of  the  candle  drawn  up  by  the 
wick,  which  fact  may  be  demonstrated  by  the  experiment 


ILLUMINATING   GAS  AND   FLAMES 


273 


shown  in  Fig.  100.  This  zone  is  also  of  relatively  low  tem- 
perature. The  next  zone  B  is  brilliantly  luminous.  Here 
partial  combustion  of  the  expanding 
gases  is  going  on.  The  ethylene  and 
other  hydrocarbon  gases  lose  their 
hydrogen,  probably  forming  at  first 
acetylene,  which  in  turn  loses  hydro- 
gen and  yields  carbon  particles  whose 
incandescence  gives  the  luminosity  to 
this  zone  of  the  flame.  Finally,  the 
outer  fringe  0  is  practically  non- 
luminous,  for  here  more  oxygen  of 
the  air  comes  into  contact  with  the 
hot  gases,  thus  causing  more  complete 
combustion.  This  latter  zone  is  the 
hottest  part  of  the  flame. 

In  the  Bunsen  flame  (Fig.  106)  the 
luminous  zone  is  absent.  We  have 
here  only  an  inner  greenish  blue  cone 
0  and  an  outer  practically  non- 
luminous  mantle  A.  In  the  inner  FlG-  107> 
cone  the  essential  processes  are  the  combustion  of  hydrogen  to 
water,  and  of  carbon  to  carbon  monoxide.  The  inner  zone 
consequently  contains  an  excess  of  reducing 
gases  and  is  termed  the  reducing  flame, 
whereas  the  outer  zone  contains  an  excess 
of  oxygen  and  is  called  the  oxidizing  flame. 
Many  of  the  metals  are  oxidized  when  intro- 
duced into  the  oxidizing  flame.  Again, 
when  oxides,  like  those  of  lead  for  instance, 
are  placed  in  the  inner  zone  they  are  re- 
duced. Blowpipe  flames  exhibit  the  same 
general  structure  as  the  Bunsen  flame. 
That  the  lower  part  of  the  inner  cone  of 
the  latter  is  relatively  low  in  temperature 
is  demonstrated  by  the  fact  that  a  match 
head  may  be  placed  in  it  for  some  time 
without  taking  fire.  The  outer  fringe  and 
the  tip  of  the  cone  near  B  are  the  hottest  parts  of  the 
flame. 


FIG.  108. 


274 


OUTLINES   OF   CHEMISTRY 


FIG.  109. 


Davy  Safety  Lamp.  —  When  a  wire  gauze  is  held  over  a 
Bunsen  burner,  and  the  gas  is  then  lighted  on  the  upper  side 
of  the  gauze,  the  flame  burns  on  that  side 
and  does  not  pass  through  the  gauze  to  the 
lower  side  (Fig.  108).  Again,  if  a  wire 
gauze  is  pressed  down  upon  a  Bunsen  flame, 
the  flame  does  not  pass  through  to  the  upper 
side  of  the  netting,  but  only  partially  con- 
sumed gas  makes  its  appearance  there  (Fig. 
109).  These  phenomena  are  due  to  the  fact 
that  the  wire  netting  lowers  the  temperature 
of  the  gases  below  the  kindling  point ;  that 
is,  the  temperature  at  which  the  gases  take 
fire  in  the  air.  If  the  gauze  should  become 
very  hot,  the  flame  will  pass  through,  of 
course. 

Upon  the  principle  that  a  wire  netting  is 
thus  able  to  intercept  a  flame  as  explained,  Sir  Humphry  Davy 
devised  the  miner's  safety  lamp  (Fig.  110).  This  consists  of 
an  oil  lamp  having  a  tight-fitting  chimney  of 
wire  gauze.  When  this  lamp  is  lighted  and 
taken  into  a  mine  where  fire  damp,  methane 
CH4,  is  present,  the  flame  is  not  communi- 
cated through  the  gauze  to  the  explosive 
mixture,  though  to  be  sure  the  latter  may 
get  into  the  chimney  through  the  gauze  and 
burn  there  or  cause  small,  harmless  explo- 
sions. These  serve  to  warn  the  miner  of  the 
presence  of  the  dangerous  gases.  The  safety 
lamp  is  consequently  very  useful ;  neverthe- 
less, explosions  do  still  occur  in  mines  because 
currents  of  air  arising  from  blasting  opera- 
tions may -blow  fine  coal  dust  into  the  lamp 
and  so  enable  the  flame  to  communicate  itself 
to  the  fire  damp  on  the  outside  of  the  gauze. 
After  such  explosions  have  occurred,  the 
carbon  dioxide  (called  choke  damp  by  the 
miners)  formed  is  dangerous  also,  because  it  does  not  support 
respiration  and  so  gives  rise  to  suffocation. 


FIG.  110. 


CHAPTER  XVII 

THERMOCHEMISTRY 

General  Remarks.  —  All  chemical  changes  are  accompanied 
by  either  an  evolution  or  an  absorption  of  heat.  In  most  of 
the  ordinary  chemical  processes  heat  is  liberated.  These  are 
consequently  called  exothermic  changes,  to  distinguish  them 
from  endothermic  changes,  or  reactions  in  which  heat  is  absorbed. 
Endothermic  changes  are  by  no  means  uncommon.  In  fact 
many  reactions  occur  with  absorption  of  heat,  particularly  at 
higher  temperatures.  It  must  be  borne  in  mind  that  physical 
changes  as  well  as  chemical  reactions  are  generally  accompanied 
by  thermal  changes.  Thus,  in  melting  ice  or  vaporizing  water 
heat  is  absorbed,  while  in  freezing  water  or  in  condensing 
va.por  heat  is  liberated.  Similarly,  whenever  a  solid  is  con- 
verted into  a  liquid,  or  a  gas  is  formed  from  a  solid  or  liquid, 
heat  is  absorbed  so  far  as  the  physical  change  is  concerned; 
and  heat  is  liberated  when  the  reverse  action  takes  place.  The 
amount  of  heat  required  to  convert  1  gram  of  a  given  solid  into 
liquid  of  the  same  temperature  is  termed  the  latent  heat  of 
fusion.  And  the  amount  of  heat  necessary  to  change  1  gram 
of  a  liquid  into  vapor  of  the  same  temperature  is  called  the 
latent  heat  of  vaporization. 

When  a  piece  of  zinc  is  dissolved  in  hydrochloric  acid,  the 
solid  zinc  disappears  and  becomes  part  of  the  liquid,  and  simul- 
taneously a  gas,  hydrogen,  is  evolved.  Both  of  these  processes, 
the  liquefaction  of  the  solid  and  the  liberation  of  the  gas,  con- 
sidered as  physical  processes,  would  proceed  with  absorption  of 
heat.  However,  the  action  of  hydrochloric  acid  on  zinc  pro- 
ceeds with  disengagement  of  heat,  which  fact  can  readily  be 
demonstrated  by  means  of  a  thermometer  placed  in  the  acid. 
It  is  therefore  evident  that  the  thermal  change  observed  is 
equal  to  the  heat  developed  by  the  chemical  interaction,  minus 
the  heat  required  for  the  liquefaction  of  the  metal  and  the 
conversion  of  the  hydrogen  into  the  gaseous  state.  It  is  at 

275 


276 


OUTLINES   OF  CHEMISTRY 


present  impossible  to  determine  just  how  much  energy  the  last- 
named  processes  represent  when  they  occur  at  room  tempera- 
ture, and  so  it  is  also  impossible  to  tell  how  much  heat  the 
actual  chemical  part  of  the  change  evolves.  All  chemical 
changes  are  similarly  accompanied  by  physical  changes  of 
some  kind.  The  thermal  effect  of  the  latter  must  be  taken 
into  consideration ;  or,  at  any  rate,  if  this  effect  cannot  be 
evaluated  and  subtracted,  as  is  frequently  the  case,  the  physi- 
cal state  of  the  substances  before  and  after  the  reaction  must 
be  mentioned. 

Calorimeters.  —  Thermal  changes  are  measured  by  means  of 
calorimeters.  A  thermometer  is  introduced  into  a  known 
weight  of  water  contained  in  a  cylindrical  dish,  the  calorimeter, 
which  is  preferably  made  of  platinum.  The  apparatus  is  so 
arranged  that  the  heat  evolved  by  the  chemical  reaction  is 
communicated  to  the  calorimeter  water.  Knowing  the  initial 

and  final  temperature  of  the 
latter,  and  multiplying  the 
weight  of  the  water  by  the  num- 
ber of  degrees  of  temperature 
change,  the  number  of  calories 
of  heat  evolved  is  obtained. 
A  large  calorie  is  the  amount 
of  heat  necessary  to  raise  1000 
grams  of  water  1  degree ;  it 
is  commonly  designated  by 
Cal.  A  small  calorie  is  0.001 
of  a  large  calorie  and  is  desig- 
nated by  cal.  In  technical 
work  in  England  and  America 
another  heat  unit  known  as 
the  British  thermal  unit  is 
frequently  used.  A  British 
thermal  unit  is  the  amount  of 
heat  required  to  raise  the 
temperature  of  1  pound  of 
water  1  degree  Fahrenheit; 
it  is  designated  by  B.  T.  U. 

During  calorimetric  measurements  care  must  be  taken  to  pre- 
vent loss  of  heat  by  radiation,  or  the  exact  amount  of  heat  lost 


THERMOCHEMISTRY 


277 


by  radiation  must  be  ascertained,  and  a  proper  correction  made 
therefor  in  the  final  result.  Figure  111  shows  an  ordinary 
calorimetric  apparatus  in  cross  section.  The  calorimeter  itself 
should  not  have  less 
than  500  cc.  capacity. 
In  Fig.  112  a  com- 
bustion calorimeter  is 
represented.  The  sub- 
stance to  be  burned  is 
placed  in  the  steel 
bomb,  which  is  lined 
with  platinum,  gold, 
or  porcelain  enamel. 
The  bomb  is  then  filled 
with  oxygen  under  20 
atmospheres  pressure 
and  finally  immersed 
in  the  water  of  the 
calorimeter.  The  ig- 
nition is  effected  by 
means  of  a  small  wire 
in  the  bomb,  heated 
by  an  electric  current. 
Thus  the  combustion 
proceeds  almost  instan- 
taneousty,  and  the  heat 
is  communicated  to  the 
calorimeter  water  and 
measured  in  the  usual 
way. 

Laws  of  Thermo- 
chemistry.—  Inasmuch 
as  energy  can  neither 
be  created  nor  de- 
stroyed, it  is  evident  that  if  no  heat  be  lost,  the  heat  evolved 
during  a  chemical  change  is  always  exactly  equal  to  the  heat  that 
is  absorbed  when  the  reaction  is  reversed.  This  law  was  pointed 
out  in  1783  by  Lavoisier  and  Laplace,  who  regarded  it  as  self- 
evident. 

In  1840  G.  H.  Hess,  professor  at  the  University  of  St.  Peters- 


FIG.  112. 


278  OUTLINES   OF   CHEMISTRY 

burg,  showed  that  the  thermal  change  accompanying  any  chemical 
reaction  depends  on  the  initial  and  final  condition  of  the  substances 
involved,  and  is  independent  of  the  intermediate  changes  that  may 
occur  during  the  reaction.  Thus  the  total  amount  of  heat  evolved 
when  a  gram  of  carbon  is  burned  to  CO2  is  the  same  whether 
the  combustion  proceeds  in  one  step,  or  whether  CO  is  first 
formed,  and  this  is  then  oxidized  to  CO2.  This  law  of  Hess 
really  follows  from  the  law  of  conservation  of  energy.  It  is 
of  great  importance  in  thermochemical  measurements,  for  it 
enables  many  determinations  to  be  made  indirectly  that  could 
not  be  carried  out  by  direct  means.  So  it  is  practically  im- 
possible to  determine  the  amount  of  heat  evolved  when  carbon 
is  burned  to  CO,  for  some  CO2  always  forms  when  this  is  at- 
tempted. But  it  is  quite  possible  to  find  the  heat  developed 
when  carbon  is  burned  to  CO2,  and  also  that  evolved  when  CO 
is  burned  to  CO2  ;  and  tlie  difference  between  these  two  experi- 
mental results  is  the  heat  evolved  when  carbon  is  burned  to 
CO.  Thus :  — 

C(solid)  +  O2(gas)  =  CO2(gas)  +          97.65  Cal. 

CO(gas)  +  O(gas)   =  CO2(gas)  +          68.       Cal. 

Therefore,  C(solid)  +  O(gas)  =  CO(gas)  +  29.65  Cal. 

These  are  typical  thermochemical  equations.  For  instance,  the 
first  one  states  that  when  12  grams  of  carbon  and  32  grams  of 
oxygen  unite,  44  grams  of  carbon  dioxide  are  formed  and  97.65 
Cal.  of  heat  are  liberated.  All  other  thermochemical  equations 
are  interpreted  similarly.  According  to  the  law  of  Lavoisier 
and  Laplace,  it  is  evident  that  if  carbon  dioxide  is  to  be  decom- 
posed into  carbon  and  oxygen,  energy  to  the  amount  of.  97.65 
Cal.  is  absorbed  during  the  process  per  every  44  grams  of  CO2. 
At  first  it  appears  peculiar  that  the  combustion  of  carbon  to 
CO  yields  but  29.65  Cal.,  whereas  the  combustion  of  CO  to 
CO2  evolves  68  Cal.  But  it  must  be  remembered  that  in  the 
first  step,  when  solid  carbon  passes  into  CO,  much  energy  is 
absorbed  in  the  process  of  vaporizing  the  carbon,  which  doubt- 
less accounts  for  the  fact  that  we  get  but  29.65  Cal.  when  car- 
bon is  burned  to  CO.  It  is  evident  that  when  furnaces  are  run 
so  that  fuel  is  but  partially  burned,  i.e.  so  that  a  considerable 
proportion  of  the  carbon  is  merely  oxidized  to  monoxide,  a 


THERMOCHEMISTRY  279 

large  proportion  of  the  energy  that  might  have  been  gained  as 
heat  is  wasted. 

The  development  of  the  subject  of  thermochemistry  is  mainly 
due  to  the  work  of  Julius  Thomsen,  who  was  professor  at  the 
University  of  Copenhagen,  and  Marcellin  Berthelot,  who  was 
professor  at  the  University  of  Paris.  In  1853  the  former  stated 
that  every  simple  or  complex  change  of  a  purely  chemical  nature  is 
accompanied  by  an  evolution  of  heat ;  and  in  1879  Berthelot  an- 
nounced that  every  change  accomplished  without  the  intervention 
of  extraneous  energy  tends  to  produce  a  substance  or  substances  in 
the  formation  of  which  the  greatest  amount  of  heat  is  disengaged. 
This  is  now  commonly  termed  Berthelot's  law  of  maximum 
work.  It  is  true  that  under  ordinary  conditions  those  re- 
actions generally  take  place  that  evolve  the  greatest  amount  of 
heat;  so  in  dissolving  metals  in  acids,  in  neutralizing  the  latter 
with  bases,  in  displacing  one  metal  by  another  in  solution,  in 
the  combustion  of  carbonaceous  substances,  etc.,  heat  is  evolved. 
Nevertheless,  Berthelot's  law,  for  which  he  contended  strongly, 
does  not  hold  rigidly ;  for,  as  already  remarked,  at  very  high 
temperatures  many  endothermic  reactions  proceed  readily. 
Furthermore,  at  ordinary  temperatures  many  changes  like  the 
interaction  of  ice  and  salt  proceed  spontaneously  with  absorp- 
tion of  heat ;  though  here  doubtless  the  amount  of  heat  required 
for  the  liquefaction  of  the  ice  and  salt  is  greater  than  that 
evolved  by  the  action  of  the  salt  on  the  ice,  whence  the  cooling 
effect  observed. 

Thermochemical  Equations.  —  As  already  stated,  it  is  cus- 
tomary to  express  the  thermal  accompaniment  of  a  chemical 
change  for  the  molecular  weight  in  grams  of  the  substances  in- 
volved. Thus  in  making  lead  iodide  from  lead  and  iodine  we 
have : — 

[Pb]  +  2[I]  =  [PblJ  +  39.8  Cal. 

indicating  that  when  207.1  grams  of  lead  and  2  x  126.92  grams 
of  iodine  unite,  460.94  grams  lead  iodide  are  formed  and  39.8 
Cal.  are  simultaneously  liberated.  The  brackets  indicate  that 
the  substances  are  in  the  solid  state.  When  liquids  come  into 
consideration  parentheses  are  used,  and  in  the  case  of  gases, 
both  brackets  and  parentheses  are  omitted.  Thus,  — 

[P]  yellow  +  3  Cl  =  (PC18)  +  76.6  Cal., 


280  OUTLINES   OF   CHEMISTRY 

means  that  when  31  grams  of  solid  yellow  phosphorus  react 
with  3  x  35.46  grams  of  gaseous  chlorine  to  form  137.38  grams 
of  liquid  phosphorus  chloride,  76.6  Cal.  are  liberated.  Thermo- 
chemical  equations  must  not  be  confounded  with  the  ordinary 
chemical  equations.  The  latter  indicate  the  direction  the  chem- 
ical change  takes  and  specify  the  nature  and  amounts  of 
the  substances  formed,  whereas  thermochemical  equations  are 
energy  equations.  For  example  the  last  equation  states  that 
the  energy  represented  in  31  grams  of  solid  phosphorus  plus  the 
energy  in  106.38  grams  of  gaseous  chlorine  is  equal  to  the  en- 
ergy in  137.38  grams  of  liquid  phosphorus  trichloride  plus 
76.6  Cal.,  at  room  temperature,  i.e.  about  18°  C.  All  other 
thermochemical  equations  are  to  be  interpreted  similarly.  It 
should  be  stated  that  the  use  of  brackets  and  parentheses  to  in- 
dicate solids  and  liquids  respectively  has  been  proposed  but  re- 
cently. It  is  a  simple  form  of  designation  which  will  probably 
be  generally  adopted. 

Different  allotropic  forms  of  an  element  contain  different 
amounts  of  energy.  Thus  when  31  grams  of  red  phosphorus 
are  converted  into  phosphorus  trichloride,  we  have :  — 

[P]  red  +  3  Cl  =  (PC18)  +  49.34  Cal.; 

therefore  from  the  last  two  equations  it  follows  that  the  conver- 
sion of  yellow  phosphorus  to  red  proceeds  with  liberation  of 
27.26  Cal.  thus:  - 

[P]  yellow  =  [P]  red  +  27.26  Cal. 

Since  thermochemical  equations  are  energy  equations,  they 
may  be  transformed  like  any  algebraic  equation,  for  instance: — 

(1)  (Hg)  +  2  Cl  =  [HgClJ  +  53.3  Cal. 

(2)  (Hg)  +  2  Cl  -  53.3  Cal.  =  [HgClJ. 

(3)  (Hg)  +  2  Cl  -  [HgClJ  =  53.3  Cal. 

(4)  (Hg)  +  2  Cl  -  [HgClJ  -  53.3  Cal.  =  zero. 

Equation  (1)  indicates  that  when  liquid  mercury  and  gaseous 
chlorine  unite  to  form  solid  mercuric  chloride,  53.3  Cal.  are  lib- 
erated. Equation  (2)  indicates  that  if  solid  mercuric  chloride 
were  transformed  into  liquid  mercury  and  gaseous  chlorine, 
53.3  Cal.  would  be  absorbed.  Equation  (3)  states  that  the 
energy  in  200  grams  mercury  plus  that  in  2  x  35.46  grams 


THERMOCHEMISTRY  281 

chlorine  is  greater  than  that  contained  in  mercuric  chloride  by 
53.3  Cal.,  and  equation  (4)  expresses  the  same  fact. 

The  total  energy  contained  in  any  substance  is  an  unknown 
quantity,  for  we  have  no  way  of  robbing  a  substance  of  all  of  its 
energy  and  measuring  the  same.  A  certain  quantity  of  energy 
may,  however,  be  obtained  from  substances ;  this  is  the  available 
or  free  energy.  It  varies  according  to  the  nature  of  the  changes 
to  which  a  substance  is  subjected.  So  by  burning  phosphorus 
in  excess  of  oxygen  more  heat  is  developed  than  by  burning  it 
in  excess  of  chlorine :  — 

2[P]  +  5  O      =  [P2O5]    +  370  Cal. 

2[P]  +  10  Cl  =  2[PC16]  +  218.4  Cal. 

It  is  clear  that  thermochemistry  can  deal  only  with  available 
energy. 

Definitions.  —  The  heat  of  solution  is  the  thermal  change  ac- 
companying the  solution  of  a  substance  in  so  much  solvent  that 
the  addition  of  more  solvent  causes  no  further  appreciable  ther- 
mal change.  The  heat  of  solution  is  commonly  stated  per  gram- 
molecule  of  dissolved  substance,  thus:  — 

[NaCl]  +  (aq)  =  (NaClaq)  -  1.3  Cal. 

indicates  that  when  1  gram-molecule  of  sodium  chloride  is  dis- 
solved in  much  water  (100  to  400  gram-molecules  of  water, 
which  is  indicated  by  aq  in  all  thermochemical  equations) 
there  is  formed  the  dilute  solution  NaClaq,  and  1.3  Cal.  are 
absorbed. 

The  heat  of  dilution  is  the  thermal  change  accompanying  the 
dilution  of  a  given  solution  with  a  definite  amount  of  pure 
solvent,  usually  so  much  that  the  addition  of  further  solvent 
does  not  cause  any  appreciable  change  of  temperature. 

The  heat  of  reaction  is  a  general  term  used  to  express  the 
thermal  change  that  accompanies  any  chemical  reaction.  The 
heat  of  formation  of  a  chemical  compound  is  the  thermal  change 
accompanying  the  formation  of  that  compound  from  the  ele- 
ments. The  terra  is  also  sometimes  used  to  indicate  the  ther- 
mal change  that  accompanies  the  formation  of  a  compound  from 
other  compounds,  or  from  elements  and  compounds.  When 
so  used,  it  is  necessary  to  specify  from  what  substances  the 
compound  whose  heat  of  formation  is  under  consideration  has 
been  formed. 


282  OUTLINES   OF   CHEMISTRY 

The  heat  of  neutralization  is  the  heat  liberated  when  an  acid 
is  neutralized  by  a  base.  The  heat  of  combustion  is  the  heat 
evolved  when  a  substance  is  completely  burned.  In  all  cases 
the  thermal  change  is  computed  per  gram-molecule. 

Thermochemical  Data.  —  These  generally  consist  of  tables  of 
heats  of  formation,  solution,  neutralization,  and  combustion. 
From  what  has  been  stated,  tables  of  this  kind  will  readily  be 
understood.  The  thermochemical  data  of  nearly  all  of  the 
important  substances  have  been  determined  by  Thomsen  and 
Berthelot. 

When  the  heat  of  formation  of  a  compound  in  solution  is 
known,  the  heat  of  formation  in  the  anhydrous  condition  may 
be  found  by  subtracting  the  heat  of  solution,  carefully  consid- 
ering the  sign  of  the  latter.  By  making  use  of  the  law  of.  Hess, 
the  heat  of  formation  of  any  compound  may  be  computed  from 
the  heat  of  any  reaction  involving  that  compound,  providing  the 
heats  of  formation  of  the  other  compounds  in  the  reaction  are 
known.  In  this  way  the  heat  of  formation  of  a  compound  from 
the  elements  may  be  found  indirectly,  even  though  it  has  not 
been  synthesized.  Thus,  let  the  heat  of  formation  of  cane 
sugar  be  required.  Its  heat  of  combustion  found  experiment- 
ally is:  — 

[C12H22On]  +  12  02  =  12  C02  +  11(H20)  +  1353  Cal. 
Again  by  experiment  it  has  been  found  that 

[C]  +  O2  =  CO2  +  97.65  Cal.,  and 
H2  +  O   =  (H2O)  +  68.4  Cal. 

It  is  clear  then  that  12  CO2  when  formed  from  12  [C]  and 
12  O2  will  liberate  12  x  97.65  Cal. ;  and  similarly  11  (H2O)  rep- 
resents a  heat  of  formation  of  11  x  68.4  Cal.  Thus,  12  CO2 
and  11  (H2O)  together  represent  a  liberation  of  12  x  97.65  + 
11  x  68.4  or  1924.2  Cal.  We  may  conceive  of  12  [C]  and 
11  H2  as  oxidized  in  one  step  to  12  CO2  and  11  (H2O)  when 
1924.2  Cal.  are  liberated  ;  or  we  may  think  of  the  operation 
as  going  on  in  two  steps  :  (1)  the  oxidation  of  12  [C]  and 
11  H2  to  sugar  (i.e.  to  [C12H22On]),  and  then  (2)  the  oxida- 
tion of  the  latter  to  12  CO2  and  11  (H2O).  Now,  since  the 
complete  oxidation  evolves  1924.2  Cal.  and  step  (2)  evolves 


THERMOCHEMISTRY  283 

1353  Cal.,  it  is  evident  that  step  (1),  which  is  the  formation 
of  sugar  from  the  elements,  proceeds  with  an  evolution  of 
1924.2  -  1353  or  571.2  Cal.  In  this  computation  97.65  Cal.  rep- 
resents the  heat  of  combustion  of  amorphous  carbon.  The 
heat  of  combustion  of  diamond  is  94.3  Cal.  If  the  latter  value 
be  employed  in  the  problem  selected,  the  heat  of  formation 
from  the  elements  will  obviously  be  571.2—12x3.35,  or 
531.3  Cal.  The  value  3.35  Cal.  clearly  represents  the  difference 
in  energy  between  amorphous  carbon  and  diamond. 

From  the  foregoing  illustration,  the  value  of  the  heats  of 
formation  of  compounds  in  computing  the  thermal  accompani- 
ments of  chemical  changes  is  evident.  The  following  tables, 
giving  the  thermochemical  data  of  a  few  of  the  most  important 
compounds,  will  serve  to  illustrate  how  such  results  are  usually 
presented:  — 


284 


OUTLINES   OF   CHEMISTRY 


TABLE   1  — HEATS   OF  FORMATION 

VALUES  ARE  EXPRESSED  IN  LARGE  CALORIES.      THE  SUBSTANCES 
NAMED  ARE  IN  THE  USUAL  STATE  AT  15°  C. 


COMPOUND 

FORMED  FROM 

GASEOUS 

LIQUID 

SOLID 

DISSOLVED 

HF 

H,F 

38.5 

45.7 



50.3 

HC1 

PI,  Cl 

22.0 





39.3 

HBr 

H,  (Br) 

86 





28.6 

HI 

H,  [I] 

-6.1 





13.1 

H2O 

H2,0 

55.3 

68.4 

69.8 



HA 

H2,02 







45.3 

H202 

(H20),  0 







-23.1 

H,S 

H2,  [S]  rhombic 

2.7 



7.3 

NH3 

N,H3 

12.0 

16.6 



20.4 

PH3 

yellow  [P],  H3 

4.3 







AsH3 

cry  st.   [As],  H3 

-44.1 







SbH3 

[Sb],  H3 

-86.8 







C2H2 

diamond  [C2],  H2 

-58.1 







C2H4 

[C2],  H4 

-14.6 







C2H6 

[C2],  H6 

23.3 







CH4 

[C],  H4 

17.3 







SiH4 

cryst.   [Si],  H4 

-6.7 







o, 

02 

-30.7 





HC10 

H,  Cl,  0 







31.65 

HC1O3 

H,  Cl,  O3 







24.0 

HC104 

H,  Cl,  04 



18.3 



38.6 

HBrO3 

H,  (Br)  03 





12.3 

HI03 

H,  [I],  03 





57.9 

55.7 

HTO4 

H,  [I],  04 



-   



47.6 

SO2 

rhombic  [S],  O2 

71.0 





78.8 

S03 

rhombic  [S]  ,  O3 





103.3 

142.5 

H2SO4 

H2,  [S],  04 



189.9 

210.9 

H2S203 

H2,  [S],  03 







141.7 

N,0 

N2,0 

-17.4 



NO 

N,  O 

-21.5 







NA 

N2,03 







6.8 

NO2 

N,02 

-7.7 







N204 

N2,04 

-2.6 







N205 

N2,06 



.  

13.1 

29.8 

HNO8 

H,  N,  03 

34.4 

41.5 

42.2 

48.8 

PA 

yellow  [P9],O5 





370.0 

406.0 

H3F04 

H3,  [P],  04 





304.1 

306.8 

As203 

[As2],  03 



154.7 

147.0 

Sb2O3 

[SbJ,  03 





166.9 



Bi203 

[Bi2],  03 



139.2 



BA 

amorph.  [B2],  O3 





272.6 



SiO2aq 

cryst.  [Si],  O2,  aq. 





179.6 



CO 

amorph.   [C]  ,  O 

29.4 







CO 

diamond  [C],  O 

26.1 







C02 

amorph.   [C],  O2 

97.65 





103.25 

THERMOCHEMISTRY 


285 


TABLE    1  —  Continued 


COMPOUND 

FORMED  FROM 

GASEOUS 

LIQUID 

SOLID 

DISSOLVED 

CO2 

diamond  [C],  O2 

94.3 





99.9 

PC13 

yellow  [P],  C13 

69.3 

76.6 





PC15 

yellow  [P],  C15 





109.2 



AsCl3 

[As],  C13 



71.7 





SbCls 

[Sb],  Cl, 





91.4 



BiC]3 

[Bi],  Cl, 





96.6 

CC14 

diamond  [C],  C14 

68.5 

75.7 





SiCl4 

cryst.  [Si],  C14 

121.8 

128.1 





SiiCl4 

[Sri],  C14 

122.2 

129.8 





(CN)2 

diamond  [C2],  N, 

-73.9 

-68.5 



67.1 

HCN 

diamond  H,"[C],~N 

-30.5 

-24.8 



24.4 

CS2 

diamond  [C],   [S2]  rhombic 

-25.4 

-19.0 





TABLE  2  — HEATS  OF   FORMATION   OF   SOME   METALLIC 
COMPOUNDS  AT   15°  C. 

VALUES  ARE  EXPRESSED  IN  LARGE  CALORIES 


COMPOUND 

FORMED  FROM 

SOLID 

DIS- 
SOLVED 

COMPOUND 

FORMED  FROM 

SOLID 

DIS- 
SOLVED 

KOH 

[K],O,  H 

103.2 

116.5 

NH4C1 

N,  H4)  Cl 

75.8 

71.9 

NaOH 

[Na],  0,  H 

101.9 

111.8 

CaCl2 

[Ca],  C12 

169.8 

187.2 

LiOH 

[Li],  O,  H 

112.3 

118.1 

ZnCl2 

[Zn],  C12 

97.2 

112.8 

NH3aq 

NH3,  w(H20) 



20.3 

A1C18 

[Al],  C18 

161.0 

237.8 

MgO 

[Mg],  0 

143.4 



FeCl2 

[Fe],  C12 

82.1 

100.0 

CaO 

[Ca],  0 

131.5 

149.6 

NiCl2 

[Ni],  C12 

74.5 

93.7 

Ca(OH)2 

[Ca],  Oa,  H2 

214.9 

217.9 

CoCl2 

[Co],  C12 

76.5 

94.8 

SrO 

[Sr],  0 

128.4 

157.7 

HgCl 

(Hg),  Cl 

31.4 



BaO 

[Ba],  0 

124.2 

158.7 

HgCl2 

(Hg),  C12 

53.3 

50.0 

BaO2 

[Ba],  02 

12.1 



AgCl 

[Ag],  Cl 

29.4 



MnO 

[Mn],  0 

90.9 



AuCl 

[Au],  Cl 

5.8 



MnO2 

[Mn],  02 

125.3 



AuCl3 

[Au],  C13 

22.8 

27.3 

FeO 

[Fe],  0 

65.7 



PtCl4 

[Pt],  C14 

59.8 

79.4 

Fe304 

[Fe3],04 

270.8 



NaBr 

[Na],(Br) 

85.8 

.83.9 

ZnO 

[Zn],  0 

84.8 



Nal 

[Na],  [I] 

69.1 

70.3 

CuO 

[Cu],  0 

37.2 



Na2C03 

[Na2],  [C],  03 

269.9 

275.4 

Cu2O 

[Cu2],  0 

40.8 



NaHC03 

[Na],H,[C],08 

227.0 

223.7 

PbO 

[Pb],  0 

50.3 



Na2S04 

[Na2],  [S],  04 

328.4 

329.0 

Pb02 

[Pb],  02 

63.4 



NaHS04 

[Na],H,[S],04 

267.8 

266.6 

HgO 

(Hg),  0 

20.1 



KN03 

[K],  N,  03 

119.5 

111.0 

KSH 

[K],[8],H 

62.3 

63.1 

KC103 

[K],  Cl,  03 

95.0 

85.0 

CaS 

[Ca],  [S] 

89.6 



KBrO3 

[K],  (Br),03 

84.1 

74.3 

SrS 

[Sr],  [S] 

97.4 



KI03 

[K],  [I],  03 

124.5 

117.4 

BaS 

[Ba],  [S] 

98.3 



KCN 

[K],  [C],  N 

29.8 

26.8 

FeS 

[Fe],  [S] 

24.0 



KMn04 

[K],  [Mn],  04 

195.0 

184.8 

CuS 

[Cu],  [S] 

10.1 



AgN03 

[Ag],  N,  08 

28.7 

23.3 

Ag2S 

[Ag2],  [S] 

3.3 



CuS04 

[Cu],  [S],  04 

1828 

198.4 

NaCl 

[Na],  Cl 

105.6 

101.2 

BaSO4 

[Ba],  [S],  04 

338.1 



286 


OUTLINES   OF   CHEMISTRY 


TABLE  3  — HEATS    OF    COMBUSTION    OF    SOME    CARBON 
COMPOUNDS 

VALUES  ARE  GIVEN  IN  LARGE  CALORIES.     SUBSTANCES  ARE  IN  THEIR 
USUAL  STATE  AT  15°  C. 


COMPOUND 

FOEMULA 

HEAT  OF 
COMBUSTION 

HEAT  OF 
FORMATION 
DIAMOND  =  [C] 

Methane 

CH4 

213.5 

16.5 

Ethane 

C2H6 

370.5 

22.1 

Propane 

C3H8 

529.2 

25.4 

Benzene 

C6H6 

787.8 

-9.1 

Methyl  alcohol 

CH3OH 

170.6 

61.7 

Ethyl  alcohol 

C2H5OH 

325.7 

69.9 

Glycerine 

C3H5(OH)3 

397.2 

161.7 

Acetic  acid 

CH3COOH 

209.4 

117.2 

Oxalic  acid 

(COOH)2 

60.2 

196.7 

Stearic  acid 

C18H36O9 

2677.8 

227.6 

Starch 

c6H10o; 

684.9 

225.9 

Dextrine 

C6H1005 

667.2 

243.6 

Cellulose 

C6H1005 

680.4 

230.4 

Cane  sugar 

C12H22On 

1353.0 

531.3 

Milk  sugar 

,  C12H22°11 

1351.4 

537.4 

Malt  sugar 

C12H22On 

1350.7 

538.1 

Dextrose 

C6H1206 

677.2 

302.6 

Lsevulose 

CCH1206 

675.9 

303.9 

Urea 

CO(NH2)2 

152.2 

77.5 

TABLE    4  — HEATS    OF    COMBUSTION    OF    VARIOUS    OTHER 
ORGANIC   SUBSTANCES 


SUBSTANCE 


HEAT  OF  COMBUSTION 
PER  1  GRAM  SUBSTANCE 


Butter 

Animal  or  vegetable  fats  and  oils,  average 

Caseine 

White  of  egg 

Egg  yolk 

Peptone  ...... 

Gluten .   . 

Muscular  tissues      ..... 

Fibrin 

Hemoglobin 


9.2  Cal. 

9.5  Cal. 

5.6  Cal. 

5.7  Cal. 
8.1  Cal. 

5.3  Cal. 
6.0  Cal. 
5.7  Cal. 
5.5  Cal. 
5.9  Cal. 


THERMOCHEMISTRY 


287 


TABLE  5  — HEATS   OF  NEUTRALIZATION 

VALUES   GIVEN  IN  LARGE   CALORIES.      (THE   SOLUTIONS   CONTAINED  1 
GRAM  EQUIVALENT  OF  ACID  OR  BASE  IN  Two  LITERS.     SOME  OF 

THE    BASES    USED    AND    SULPHATES    FORMED    ARE    INSOLUBLE.) 


BASES 

HClaq 

HN03  aq 

CH3COOH  aq 

|H2S04aq 

HCNaq 

KOHaq 

13.7 

13.8 

13.3 

15.7 

3.0 

NaOH  aq 

13.7 

13.7 

13.3 

15.85 

2.9 

NH4OH  aq 

12.45 

12.5 

12.0 

14.5 

1.3 

£  Ca(OH)2  aq 

14.0 

13.9 

13.4 

15.6 

3.2 

1  Sr(OH)2  aq 

14.0 

13.9 

13.3 

15.4 

3.1 

£  Ba(OH)0  aq 

13.85 

13.9 

13.4 

18.4 

3.2 

Mg(OH)2  aq 

13.8 

13.8 



15.6 



£  Fe(OH)2  aq 

10.7 



9.9 

12.5 



Ni(OH)2  aq 

11.3 





13.1 



$  Co(OH)2  aq 

10.6 





13.3 



JZn(OH)saq 

9.8 

9.8 

8.9 

11.7 



Cu(OH)2  aq 

7.5 

7.5 

6.2 

9.2 



Uses  of  Thermochemical  Data.  —  From  Table  1  it  appears  that 
the  heat  of  formation  of  the  hydrohalogens  diminishes  as  the 
atomic  weight  of  the  halogens  increases,  hydrfodic  acid  even 
having  a  negative  heat  of  formation.  We  have  seen  that  the 
stability  of  these  compounds  diminishes  in  the  same  way,  hydro- 
fluoric acid  being  the  stablest  and  hydriodic  acid  the  least  stable. 
On  the  other  hand,  it  will  be  recalled  that  iodic  and  periodic 
acids  are  more  stable  than  chloric,  bromic,  and  perchloric  acids, 
and,  indeed,  Table  1  shows  that  the  heats  of  formation  of  iodic 
and  periodic  acids  are  higher  than  those  of  the  corresponding 
compounds  of  the  other  halogens.  Water  is  a  much  stabler 
compound  than  hydrogen  sulphide,  as  is  borne  out  by  the  great 
difference  in  their  heats  of  formation.  Ammonia,  phosphine, 
arsine,  and  stibine  diminish  in  stability  in  the  order  named, 
which  is  precisely  what  one  would  expect  from  their  heats  of 
formation,  which  also  diminish  in  the  same  order.  Marsh  gas 
and  ethane,  it  will  be  observed,  have  positive  heats  of  forma- 
tion, whereas  the  unsaturated  compounds  ethylene  and  acety- 
lene are  formed  with  absorption  of  heat.  The  formation  of 
ozone  from  oxygen  takes  place  with  absorption  of  much  energy, 
as  the  negative  heat  of  formation  of  ozone  indicates.  From 
these  illustrations  and  from  others  with  which  Tables  1  and  2 


288  OUTLINES   OF   CHEMISTRY 

are  replete,  it  appears  that  thermochemical  data  offer  a  means 
of  comparing  the  relative  stability  of  compounds. 

Both  Thomsen  and  Berthelot  had  hoped  that  thermochemical 
data  would  offer  a  means  of  exact  measurement  of  chemical 
attractions,  but  this  has  not  been  realized.  Thermochemical  data 
are  complicated  by  the  fact  that  they  also  represent  the  energy 
concomitants  of  physical  changes  which  invariably  accompany 
chemical  reactions,  and  which  cannot  be  evaluated,  as  already 
explained.  Moreover,  it  must  be  borne  in  mind  that,  in  speak- 
ing of  the  stability  of  a  substance,  it  is  really  necessary  to  specify 
toward  what  agencies  such  stability  is  being  considered.  Thus, 
a  substance  A  might  be  much  stabler  than  another  substance  B 
towards  the  decomposing  action  of  heat,  whereas  towards  the 
action  of  electricity,  light,  or  the  inroads  of  various  reagents, 
A  might  be  less  stable  than  B.  So,  for  instance,  carbon  tetra- 
chloride,  with  its  heat  of  formation  +68.5,  ought  to  be  less 
stable  than  silicon  tetrachloride,  whose  heat  of  formation  is 
+  121.8.  While  this  is  substantiated  by  the  fact  that  silicon 
tetrachloride  may  readily  be  obtained  by  passing  chlorine  over 
hot  silicon,  whereas  carbon  tetrachloride  cannot  be  similarly 
obtained,  it  must  also  be  borne  in  mind  that  when  treated  with 
water,  silicon  tetrachloride  is  at  once  decomposed  into  hydro- 
chloric and  silicic  acids,  whereas  carbon  tetrachloride  remains 
unchanged  under  the  same  treatment.  However,  here  the  fact 
that  the  heat  of  formation  of  silicic  acid  by  far  exceeds  that  of 
carbonic  acid  no  doubt  is  a  determining  factor.  By  means  of 
the  electric  current  neither  SiCl4  nor  CC14  can  be  decomposed, 
whereas  common  salt,  which  per  gram  equivalent  has  over  five 
times  as  high  a  heat  of  formation  as  carbon  tetrachloride,  is 
nevertheless  easily  decomposed  by  electrolysis  (which  see). 
Thus  it  is  clear  that  great  care  must  be  exercised  in  using  thermo- 
chemical data  in  arguing  as  to  the  relative  stability  of  compounds. 

The  value  of  fuels  depends  upon  their  heat-giving  power; 
that  is,  their  heat  of  combustion.  And  so  it  is  clear  that  the 
heats  of  combustion  of  wood,  coal,  and  various  liquid  and  gas- 
eous fuels  is  of  utmost  practical  importance.  In  the  animal 
body  the  foods  consumed  are  digested,  assimilated,  and  finally 
slowly  oxidized  and  eliminated  in  the  form  of  carbon  dioxide 
and  water  in  the  case  of  carbohydrates  and  fats,  and  in  the  form 
of  carbon  dioxide,  water,  urea,  and  other  nitrogenous  products 


THERMOCHEMISTRY  289 

in  the  case  of  nitrogenous  foods.  Therefore  the  heats  of  com- 
bustion of  foodstuffs  have  sometimes  been  considered  in  deter- 
mining the  value  of  various  foods.  In  such  a  procedure  great 
care  must  again  be  exercised ;  for  foods  that  have  nearly  the 
same  heat  of  combustion  are  frequently  of  quite  different  value, 
because  they  are  not  all  digested  and  assimilated  with  equal 
readiness.  Compare,  for  example,  the  heats  of  combustion  of 
starch  and  cellulose  in  Table  3  ;  the  values  are  nearly  the  same, 
and  yet  the  food  value  of  the  substances  to  an  animal  is  very 
different. 

An  inspection  of  the  heats  of  combustion  in  Table  3  shows 
that  analogous  substances  of  the  same  carbon  and  hydrogen 
content  have  approximately  the  same  heats  of  combustion,  in 
spite  of  their  differences  in  structure.  .  Nevertheless,  differences 
in  structure  do  yield  corresponding  differences  in  heats  of  com- 
bustion. This  matter  has  been  studied  in  some  detail,  espe- 
cially by  Stohmann.  Adjacent  members  of  homologous  series 
on  the  average  show  a  difference  of  about  158  Cal.  for  CH2. 
The  heat  of  combustion  of  carbon  compounds  is  approximately 
an  additive  property.  In  Table  4  are  given  the  heats  of  com- 
bustion of  a  few  additional  important  substances. 

It  will  be  observed  that  the  heats  of  neutralization  of  differ- 
ent bases  by  different  acids,  Table  5,  are  approximately  the 
same  in  the  case  of  the  strong  bases  and  strong  acids.  This 
will  be  discussed  in  connection  with  the  subject  of  electrolytic 
dissociation.  In  general,  Table  5  shows  that  when  a  given  acid 
is  neutralized,  the  heat  thus  developed  by  bases  that  are  known 
to  be  closely  related  chemically  is  approximately  the  same. 
So  when  hydrochloric  acid  is  neutralized  by  sodium  or  potassium 
hydroxide,  the  heat  of  neutralization  is  13.7  Cal.  When  the 
same  acid  is  neutralized  by  ferrous,  cobaltous,  or  nickelous 
hydroxide,  the  heat  developed  is  about  10.8  Cal. 


CHAPTER   XVIII 

SILICON  AND  BORON  AND  THEIR  IMPORTANT  COMPOUNDS 

Occurrence,  Preparation,  and  Properties  of  Silicon.  —  Next  to 
oxygen,  silicon  is  the  most  abundant  element  found  in  the 
earth's  crust,  constituting  more  th#n  one  fourth  of  the  latter. 
Silicon  does  not  occur  in  the  free  state.  It  is  always  found  in 
combination  with  other  elements,  especially  with  oxygen  as 
silica,  and  with  oxygen  and  various  metals  as  silicates.  Quartz, 
quartzite,  flint,  and  the  white  sands  of  the  seashore  and  the 
deserts  are  nearly  pure  silicon  dioxide;  whereas  clays  are  largely 
composed  of  silicates. 

Silicon  was  first  prepared  in  pure  form  in  1823  by  Berzelius, 
who  heated  potassium  silicofluoride  with  metallic  potassium  :  —  • 

K2SiF6  +  4  K  =  6  KF  +  Si. 

The  element  may  also  be  obtained  by  heating  finely  powdered 
quartz  sand  with  magnesium  powder  :  — 

SiO2  +  2  Mg  =  2  MgO  +  Si. 

In  this  case  magnesium  silicide  Mg2Si  is  generally  also  formed  ; 
but  by  means  of  hydrochloric  acid  the  silicon  can  readily  be 
freed  from  this  silicide  and  also  from  the  oxide  of  magnesium. 
Silicon  may  also  be  obtained  by  heating  sodium  or  aluminum 
in  a  current  of  silicon  tetrachloride  vapor,  thus  :  — 


4Na  =  4NaCl+Si. 
3  SiCl4  +  4  Al  =  4  A1C13+  3  Si. 

On  a  large  scale,  silicon  is  now  manufactured  at  Niagara  Falls 
by  heating  together  quartz  sand  and  coke  in  the  electric  furnace, 


thus  :  — 


SiO2  +  2  0  =  2  CO  +  Si. 


Silicon  is  run  out  of  the  electric  furnaces  into  molds.  It  thus 
forms  "pigs"  that  weigh  from  600  to  800  pounds.  The 
material  varies  in  purity  from  90  to  97  per  cent,  though  silicon 

290 


SILICON  AND   BORON  291 

over  99  per  cent  pure  has  thus  been  prepared.  Silicon  is  sold 
in  car  lots  at  about  $ 120  per  ton.  It  is  mainly  used  in  the  steel 
industry  as  a  reducing  agent.  In  1908,  500  tons  of  90  per 
cent  silicon  were  used  in  manufacturing  steel.  It  is  very  likely 
that  silicon  will  be  used  for  many  other  purposes  in  the  near 
future. 

Silicon  is  either  crystalline,  or  an  amorphous  brown  powder. 
In  the  latter  form  it  is  commonly  obtained  by  the  first  three 
methods  above  described.  Amorphous  silicon  burns  when 
highly  heated  in  the  air,  the  product  being  silicon  dioxide  SiO2. 
Since  the  latter  is  practically  non-volatile,  its  accumulation 
hinders  the  securing  of  complete  oxidation  of  all  the  silicon. 
Under  a  layer  of  common  salt,  amorphous  silicon  may  be  melted, 
and  on  cooling  it  becomes  crystalline.  Silicon  may  also  be 
obtained  in  crystalline  form  by  dissolving  amorphous  silicon  in 
molten  zinc  ;  on  cooling,  silicon  separates  out  in  form  of  crystals. 
The  zinc  may  be  removed  with  hydrochloric  acid.  Silicon 
crystallizes  in  the  isometric  system,  forming  dark  gray  shining 
plates  or  rods,  which  in  reality  consist  of  octahedra  that  have 
grown  together  so  as  to  form  twin  crystals.  The  specific  gravity 
of  silicon  is  2.49.  It  is  so  hard  that  it  will  scratch  glass.  The 
crystalline  variety  conducts  electricity,  though  rather  poorly. 
The  amorphous  powder  is  a  non-conductor.  Like  graphite, 
crystalline  silicon  is  hard  to  oxidize  by  heating  it  in  the  air  or 
in  oxygen.  Hydrofluoric  acid  attacks  it  but  slowly ;  nitric 
and  hydrofluoric  acids  act  on  it  more  rapidly.  Fluorine  reacts 
with  silicon  even  at  ordinary  temperatures  with  evolution  of 
light  and  heat :  — 

Si  +  4F=SiF4. 

Hot  solutions  of  caustic  potash  dissolve  silicon :  — 
2  KOH  +  H20  +  Si  =  K2Si03  +  2  H2. 

The  atomic  weight  of  silicon  is  28.3.  It  has  been  determined 
by  analyzing  its  compounds  with  the  halogens.  Silicon  is  quad- 
rivalent in  all  of  its  compounds,  the  formulae  of  which  conse- 
quently are  analogous  to  those  of  the  compounds  of  carbon. 
Indeed,  silicon  and  carbon  bear  many  resemblances  in  their 
chemical  behavior,  and  while  carbon  is  exceedingly  important 
in  the  organic  world,  silicon  plays  a  similar  role  in  the  inorganic 
realm. 


292 


OUTLINES  OF   CHEMISTRY 


Silicon  Dioxide,  Silica.  —  This  is  by  far  the  most  important 
compound  of  silicon.  Its  formula  is  SiO2.  It  is  silicic  acid 
anhydride.  In  the  form  of  quartzite,  it  often  forms  mountains. 
It  is  the  chief  constituent  of  sandstones,  and  sand  is  largely  silica. 
In  crystalline  form  it  occurs  as  quartz  and  amethyst,  and  also, 
though  rarely,  as  tridymite.  In  the  amorphous  form  it  is  found 
as  agate,  opal,  flint,  carnelian,  and  chalcedony,  which  frequently 
contain  water  in  combination.  Pure  silicon  dioxide  is  colorless, 
but  many  of  the  varieties  found  in  nature  are  colored  by  vari- 
ous impurities.  Thus  smoky  quartz  is  discolored  with  organic 
matter,  rose  quartz  with  manganese,  carnelian  with  oxide  of 
iron,  etc.  Quartz  crystallizes  in  the  hexagonal  system.  Its 

crystals  occur  in  two  forms 
that  are  non-superposable 
(Fig.  113);  that  is,  they 
are  to  each  other  as  the 
right  hand  is  to  the  left. 
These  crystals  rotate  the 
plane  of  polarized  light 
passed  through  them  paral- 
lel to  the  main  axis.  The 
Fia.  113.  , 

degree  of  rotation  is  propor- 
tional to  the  thickness  of  the  layer  traversed,  and  the  deviation 
is  either  dextro  or  leevo  according  to  the  crystal  used.  This 
property  makes  quartz  useful  in  certain  kinds  of  optical  in- 
struments, particularly  in  certain  types  of  polariscopes.  Tri- 
dymite also  crystallizes  in  the  hex- 
agonal system.  It  usually  occurs  in 
prismatic  plates  (Fig.  114). 

Quartz  is  brittle  and  very  hard. 
It  is  consequently  used  as  an  abra- 
sive material  in  grinding  glass,  metals,  etc.  Glued  on  paper, 
it  forms  sandpaper.  Its  specific  gravity  is  2.6.  It  requires 
the  temperature  of  the  oxyhydrogen  flame  to  melt  quartz. 
When  thus  heated,  it  forms  a  viscous  liquid  that  can  be  drawn 
out  and  worked  like  glass.  In  the  electric  furnace,  the 
liquid  may  be  boiled  and  evaporated.  In  recent  years  flasks, 
crucibles,  evaporating  dishes,  and  other  utensils  have  been  made 
of  quartz  glass.  These  have  the  great  advantage  that  they  will 
not  break  when  subjected  to  sudden  and  very  great  differences 


FIG.  114. 


SILICON  AND   BORON  293 

of  temperature.  This  is  due  to  the  'fa.ct  that  quartz  changes  its 
volume  but  very  slightly  with  alterations  of  temperature.  The 
coefficient  of  expansion  of  quartz  between  0°  and  1000°  is  only 
0.0000007  on  the  average,  being  less  than  that  of  any  other 
known  substance.  A  white-hot  quartz  crucible  may  be 
quenched  in  cold  water  without  injuring  the  dish. 

Silica  constitutes  about  40  per  cent  of  the  ash  of  the  feathers 
of  birds.  It  is  also  found  in  egg  albumin,  in  the  hair  of  ani- 
mals, and 'in  various  crustaceans.  Diatomic  or  infusorial  earth 
consists  of  the  siliceous  remains  of  minute  organisms  called 
diatoms  or  infusoria.  The  stalks  of  grasses,  cereals,  field 
horsetails,  bamboo,  and  other  canes  contain  notable  amounts  of 
silica,  which  is  in  combination  with  other  elements  and  aids  in 
giving  the  stalks  stability.  Sometimes  over  half  of  the  ash  of 
these  stalks  consists  of  silica. 

Besides  being  used  as  an  abrasive  material,  silica  is  employed 
in  the  manufacture  of  glass  and  in  making  mortar,  cement,  and 
porcelain. 

Silicic  Acids.  —  Silicon  dioxide  is  the  anhydride  of  a  series  of 
silicic  acids.  These  may  all  be  considered  as  composed  of  silica 
and  water  in  various  proportions.  They  may  all  be  referred  to 
orthosilicic  acid  Si(OH)4,  which  is  well  known  in  the  form  of 
salts,  though  it  has  not  been  prepared  in  the  pure  state.  The 
acid  is  probably  present  in  the  gelatinous  precipitate  formed 
when  silicon  tetrachloride  or  tetrabromide  is  treated  with 
water :  —  v 

SiCl4  +  4  H20  =  4  HC1  +  Si(OH)4. 

By  losing  a  molecule  of  water,  orthosilicic  acid  passes  over 
into  metasilicic  acid  H2SiO3.  From  two  molecules  of  ortho- 
silicic  acid  by  elimination  of  one,  two,  and  three  molecules  of 
water  the  following  acids,  commonly  known  as  disilicic  acids, 
are  formed:  — 

H6Si207,   H4Si206,  H2Si206. 

From  three  molecules  of  the  ortho  acid  by  loss  of  two  and  four 
molecules  of  water  the  trisilicic  acids  H8Si3O10  and  H4Si3O8  are 
formed.  None  of  these  polysilicic  acids  have  been  isolated. 
Their  existence  is  simply  vouchsafed  by  the  fact  that  salts  of 
these  acids  occur  in  nature,  or  have  been  made  in  the  labora- 
tory. The  mineral  olivine  Mg2SiO4  (Fig.  70)  is  a  salt  of 


294 


OUTLINES  OF  CHEMISTRY 


H4SiO4 ;  sodium  silicate,  or  water  glass,  Na2SiO3  is  a  salt  of 
H2SiO3 ;  serpentine  Mg3Si2O7  is  a  salt  of  H6Si2O7  ;  and  the  feld- 
spars, orthoclase  AlKSi3O8  and  albite  AlNaSi3O8,  are  salts  of 
H4Si308. 

When  silica  is  fused  with  sodium  carbonate,  sodium  silicate 
is  formed :  — 

Na2CO3  +  SiO2  =  Na2SiO3  +  CO2. 

Sodium  silicate  is  soluble  in  water  and  is  known  as  water  glass, 
as  is  also  the  silicate  of  potassium  K2SiO3,  which  may  be  made 
similarly.  The  silicates  of  metals  other  than  the  alkalies  are  very 
slightly  soluble  in  water.  On  treating  a  solution  of  sodium  or 
potassium  silicate  with  a  mineral  acid,  silicic  acid  is  set  free  :  — 

Na2SiO3  +  2  HC1  =  2  NaCl  +  H2SiO3. 

If  the  solution  is  concentrated,  the  silicic  acid  is  precipitated 
in  the  form  of  a  jelly.  If  dilute  solutions  are  used  and  the  water 
glass  is  poured  into  an  excess  of  hydrochloric  acid,  no  precipi- 
tate forms.  From  this  clear  solution,  the  sodium  chloride  and 
excess  of  hydrochloric  acid  may  be  removed  by  dialysis. 
The  apparatus  required  for  the  purpose  is  called  a  dialyser,  a 

common  form  of  which 
is  shown  in  Fig.  115. 
A  parchment  paper  or 
animal  bladder  is  se- 
curely tied  over  one  end 
of  a  cylinder  into  which 
the  solution  is  poured. 
The  whole  is  then  im- 
mersed in  a  larger  outer 
dish  of  water  as  shown 
in  the  figure.  The  so- 
dium chloride  and  hy- 
drochloric acid  pass 
through  the  septum  into 
the  outer  liquid,  while 
the  silicic  acid  remains  behind  in  the  inner  vessel.  By  renewing 
the  water  in  the  outer  dish  from  time  to  time,  practically  all  of 
the  chlorides  can  be  removed  from  the  inner  vessel,  which  then 
contains  only  a  solution  of  silicic  acid.  This  may  be  concentrated 


FIG.  115. 


SILICON  AND   BORON  295 

by  careful  evaporation  to  about  10  per  cent,  if  not  quite  all  of  the 
chlorides  have  been  removed,  or  to  about  1  per  cent  if  practically 
all  the  chlorides  have  been  taken  out.  If  attempts  are  made 
to  concentrate  to  a  greater  extent  or  to  preserve  the  solution 
for  a  long  time,  the  silicic  acid  largely  separates  out  in  form  of 
a  gelatinous  mass,  which  is  termed  a  hydrogel,  the  clear  solu- 
tion from  which  the  latter  has  been  formed  being  termed  a 
hydrosol.  The  solution  of  silicic  acid  obtained  by  dialysis  as 
described  is  also  commonly  called  a  colloidal  solution.  This 
term  was  introduced  by  Thomas  Graham  to  denote  solutions  of 
non-crystalline  bodies  which  do  not  pass  through  membranes 
used  in  dialysis  experiments.  Thus  Graham  found  that,  like 
silicic  acid,  substances  such  as  glue,  gums,  albumin,  ferric 
hydroxide,  etc.,  which  are  non-crystalline,  do  not  pass  through 
parchment  or  animal  membranes  as  readily  as  crystalline  sub- 
stances. He  consequently  made  two  classes  of  substances  : 
colloids,  which  do  not  pass  through  membranes  on  dialysis  ;  and 
crystalloids,  which  do  make  their  way  through  such  septa 
readily.  Though  this  distinction  is  still  frequently  made,  it 
really  cannot  be  held  in  the  light  of  more  recent  experiments ; 
for  it  is  quite  possible  to  separate  crystalline  substances  from 
each  other  by  this  process.  It  is  even  possible  to  effect  the 
separation  of  crystalline  from  non-crystalline  substances  by 
having  the  latter  pass  through  the  septum  and  the  crystal- 
line substances  remain  behind.  It  all  depends  upon  the  nature 
of  the  septum  chosen  and  the  character  of  the  substances  under 
consideration.  So  when  cane  sugar  and  camphor,  both  crystal- 
line substances,  are  dissolved  in  pyridine,  and  the  solution  is 
separated  from  pure  pyridine  by  means  of  a  vulcanized  caout- 
chouc membrane,  such  as  the  dentists  use  as  "rubber  dam," 
camphor  passes  through  and  sugar  remains  behind.  Again, 
when  copper  oleate,  a  non-crystalline  substance,  and  cane  sugar 
together  in  pyridine  solution  are  similarly  subjected  to  dialy- 
sis, the  copper  oleate  passes  through  the  rubber  membrane, 
and  the  crystalline  sugar  remains  behind.  Finally,  if  -to  a  solu- 
tion of  collodion  in  alcohol  and  ether,  copper  oleate  is  added  and 
this  solution  is  then  (by  means  of  a  rubber  membrane)  separated 
from  a  mixture  of  alcohol  and  ether  such  as  is  used  in  making 
up  the  collodion  copper  oleate  solution,  the  copper  oleate  passes 
through  the  septum  and  the  nitrocellulose  remains  behind, 


296  OUTLINES   OF   CHEMISTRY 

As  both  copper  oleate  and  nitrocellulose  are  non-crystalline 
in  character,  we  have  here  a  case  of  the  separation  of  two 
non-crystalloids,  that  is,  in  Graham's  language,  of  two  colloids, 
by  dialysis. 

On  drying  gelatinous  silicic  acid,  it  loses  water  and  finally 
forms  a  white  amorphous  powder  which  must  be  heated  in 
the  blast  to  expel  all  traces  of  moisture. 

Action  of  Water  on  Silicates. — Silicates  are  difficultly  soluble 
in  water,  yet  the  earth's  crust  is  continually  being  worn  away 
by  the  solvent  action  of  rain  water  upon  the  siliceous  geological 
deposits.  Rocks  like  granite,  gneiss,  schists,  shales,  and  slates 
are  continually  being  washed  away  by  the  solvent  action  of 
water,  slight  though  it  be.  Thus  a  gradual  leveling  process 
is  going  on  which  is  aided  by  the  action  of  wind  and  the 
disintegrating  effects  of  alternate  freezing  and  thawing.  So 
the  silicates  are  gradually  dissolved,  and  the  more  resistant 
quartz  grains  remain  behind  as  sand.  This,  however,  finds 
its  way  into  the  sea  and  other  depressions  filled  with  water, 
where  the  sand  grains  are  frequently  gradually  cemented 
together  with  calcium  carbonate  or  oxides  of  iron,  thus 
forming  so-called  sandstones.  As  silicic  acid  is  a  very 
weak  acid,  which  is  evident  from  the  fact  that  its  solutions 
neither  react  toward  litmus  nor  have  any  taste,  we  should 
expect  solutions  of  silicates  to  contain  these  salts,  largely 
in  a  state  of  hydrolytic  decomposition  ;  and  such  is  actually 
the  case. 

Decomposition  of  Silicates  in  the  Laboratory.  —  This  is  effected 
by  fusing  the  pulverized  silicate  with  sodium  carbonate  or  a 
mixture  of  this  salt  and  potassium  carbonate.  In  this  way 
sodium  silicate  is  formed  which  is  soluble  in  water.  The 
other  bases  present  may  generally  be  readily  dissolved  with 
the  aid  of  hydrochloric  acid.  Silicates  may  also  be  decomposed 
with  hydrofluoric  acid,  or  with  this  and  hydrochloric  or  sul- 
phuric acid.  Thus  the  silicon  is  volatilized  in  form  of  SiF4,  and 
the  bases  remain  as  chlorides  or  sulphates.  Silicates  may  further 
be  decomposed  by  heating  with  calcium  carbonate  and  ammo- 
nium chloride  and  then  extracting  the  mass  with  water.  In 
the  latter  process  calcium  silicate  is  formed,  and  the  bases  are 
converted  into  chlorides. 

Hydrogen    Silicide.  —  When    magnesium    silicide   Mg2Si   is 


SILICON  AND   BORON  297 

treated  with  hydrochloric  acid,  hydrogen  silicide  or  silico- 
methane  SiH4  is  formed  :  — 

Mg2Si  +  4  HC1  =  2  MgCl2  +  SiH4. 

The  colorless  gas  so  obtained  always  contains  hydrogen  and 
some  silicoethane  Si2H6.  Pure  SiH4  does  not  take  fire  in  the 
air  except  under  diminished  pressure.  Silicoethane,  however, 
ignites  spontaneously  on  exposure  to  the  air,  and  it  is  this  gas 
whose  presence  causes  SiH4  to  burn  in  contact  with  air  at 
ordinary  pressure.  Silicon  tetrahydride  may  be  liquefied  at 
— 11°  under  a  pressure  of  50  atmospheres.  In  chlorine  gas 
SiH4  takes  fire.  On  burning  silicon  hydride  in  the  air,  the 
products  formed  are  water  and  silica ;  the  latter  forms  white 
smoke.  Silicoethane  Si2H6  boils  at  +  52°. 

Compounds  of  Silicon  with  the  Halogens.  — Silicon  tetrafluoride 
SiF4  is  formed  by  treating  silicon  with  fluorine,  or  more  readily 
by  treating  silica  with  hydrofluoric  acid  or  a  mixture  of  fluor- 
,spar  CaF2  and  sulphuric  acid,  thus  :  — 

CaF2  +  H2SO4  =  2  HF  +  CaSO4,  and 

SiO2  +  4  HF  =  2  H2O  +  SiF4 ;  or 
2  CaF2  +  Si02  +  2  H2SO4  =  2  CaSO4  +  2  H2O  +  SiF4. 

Silicon  tetrafluoride  is  a  colorless  gas  of  ve/y  pungent  odor. 
It  boils  at  -65°,  and  the  solid  melts  at  --770.  The  tetra- 
fluoride is  always  formed  when  hydrofluoric  jacid  acts  on  sili- 
cates, and  it  is  consequently  produced  when; that  acid  is  used 
in  etching  glass.  / 

Water  decomposes  silicon  tetrafluoride  :  — 

3  H2O  +  3  SiF4  =  H2Si03  +  I  H2SiF6. 

I 

The  silicic  acid  formed  separates  out  as  i*  gelatinous  precipitate, 

while  the  hydrofluosilicic  acid  H2SiF6  remains  in  solution. 
The  latter  may  be  concentrated  to  som£  extent  by  evaporation. 
The  concentration  must  be  carried  on  ima  platinum  dish,  because 
hydrogen  fluoride  is  formed  during  the  process,  and  so  glass  or 
porcelain  dishes  would  be  attacked.  Pure  H2SiF6  is  not  known, 
for  on  attempting  to  concentrate  its  solutions  beyond  a  certain 
point  the  acid  breaks  up,  yielding  hydrogen  fluoride  and  silicon 
tetrafluoride.  In  making  fluosilicic  acid  the  silicon  tetra- 


298 


OUTLINES   OF   CHEMISTRY 


fluoride  generated  in  a  flask  by  the  reaction  above  mentioned  is 
conducted  into  water  by  means  of  a  tube  whose  lower  end  dips  in 

mercury  (Fig.  116),  so  that 
the  gelatinous  silicic  acid 
formed  will  not  stop  the  end  of 
the  tube.  As  the  gas  rises  from 
the  mercury,  clouds  of  silicic 
acid  are  formed  in  the  water. 
Hydrofluosilicic  acid  is 
a  strong  acid.  It  readily 
decomposes  carbonates  and  hy- 
droxides of  the  metals,  form- 
ing the  fluosilicates.  The 
latter  are  decomposed  by 
heat,  yielding  fluorides  of  the 
metals  and  silicon  tetra- 
fluoride.  The  silicofluorides 
are  commonly  soluble  in  water, 
insoluble,  and  the  potassium  salt  is 


FIG.  116. 


The    barium    salt    is 
sparingly  soluble. 

Silicon  tetrachloride  SiCl4  is  formed  by  heating  silicon  in  a 
current  of  chlorine,  or  by  passing  chlorine  over  a  heated  mix- 
ture of  carbon  ar^d  silica,  thus :  — 

Si  +  2  C12  =  SiCl4,  or 
SiOj  +  2  C  +  C14  =  2  CO  +  SiCl4. 

The  product  is1  a  liquid  of  pungent  odor.  It  boils  at  59°,  has 
a  specific  gravity  of  1.52  at  0°,  and  solidifies  at  —  89°.  Water  at 
once  decomposes  it  t  — 

SiCl4  -P  4  H20  =  4  HC1  +  H4Si04. 

On  heating  silicon  jn  a  current  of  hydrochloric  acid  gas, 
silicon  chloroform  SiHCl3  may  be  obtained.  This  boils  at  34°, 
has  a  specific  gravity  of  1.3,  and,  like  silicon  tetrachloride,  it  is 
at  once  decomposed  by  ?vater. 

Bromine  and  iodine  compounds  of  silicon,  analogous  to  the 
chlorine  compounds,  have  been  prepared  by  similar  methods. 
Silicon  tetrabrpmide  boils  at  153°  and  melts  at  -  12°,  silicon 
tetraiodide  SiI4  forms  octahedra  that  melt  at  120°  and  boil  at 
290°. 


SILICON  AND   BORON  299 

Esters  of  Silicic  Acid.  —  Methyl  silicate  (CH3)4SiO4,  boiling 
at  121°,  and  ethyl  silicate  (C2H5)4SiO4,  boiling  at  165°,  are  also 
known.  They  are  formed  by  the  action  of  alcohols  on  silicon 
tetrachloride,  thus  :  — 

SiCl4  +  4  CH3OH  =  4  HC1  +  (CH3)4SiO4. 

Water  decomposes  the  esters  to  alcohol  and  silicic  acid. 

Silicon  Carbide,  Carborundum,  SiC.  —  This  substance  is 
formed  in  the  electric  furnace  by  heating  together  silica  or 
quartz  sand,  carbon,  and  common  salt  to  about  3500°.  The 
following  reaction  occurs  :  — 


Silicon  carbide  forms  hexagonal  plates  that  commonly  have  a 
dark  greenish  blue  color.  The  substance  is  not  attacked  by 
acids;  not  even  hydrofluoric  acid  makes  inroads  upon  it.  It 
may  readily  be  decomposed,  however,  by  fusion  with  caustic 
alkalies.  Carborundum  has  a  specific  gravity  of  3.2,  and  is 
extremely  hard,  being  next  to  the  diamond  in  hardness.  It  is 
consequently  used  as  an  abrasive  material.  Grinding  wheels, 
whetstones,  etc.,  made  of  carborundum  are  in  common  use. 

Titanium,  Zirconium,  and  Thorium.  —  These  are  quadrivalent 
metallic  elements  whose  compounds  are  analogous  to  those  of 
silicon.  The  elements  are  steel-gray,  brittle  metals. 

Titanium  (Ti  —  48.1)  is  found  in  nature  as  the  dioxide  TiO2, 
in  form  of  rutile,  brookite,  and  anatase.  The  element  is  widely 
distributed,  but  occurs  nowhere  in  large  quantities.  It  is  also 
met  in  titaniferous  iron  ores,  which  are  in  the  main  ferrous 
titanate  FeTiO3.  It  also  occurs  together  with  zircon  in  certain 
silicates. 

Zirconium  (Zr  —  90.6)  is  found  chiefly  in  the  mineral  zircon, 
which  forms  tetragonal  crystals  of  the  composition  ZrSiO4, 
from  which  Klaproth  prepared  the  dioxide  ZrO2  in  1789. 
Moissan  prepared  the  metal  by  heating  the  oxide  with  carbon 
in  the  electric  furnace. 

Though  in  its  compounds  cerium  is  also  more  frequently 
quadrivalent,  it  will  nevertheless  be  discussed  in  connection 
with  lanthanum  and  other  rare-earth  elements  (which  see). 

Thorium  (Th  —  232.4)  was  found  in  thorite  ThSiO4  •  2  H2O 
by  Berzelius,  in  1828.  Thorium  salts  are  now  prepared  from 


300  OUTLINES   OF   CHEMISTRY 

monazite  found  in  North  Carolina.  Welsbach  light  mantles 
consist  of  99  per  cent  thoria  ThO2  and  1  per  cent  ceria  CeO2 
(see  under  cerium).  Thorium  compounds  are  radio-active 
(see  radium). 

Occurrence,  Preparation,  and  Properties  of  Boron.  —  This  ele- 
ment occurs  in  nature  in  the  form  of  boric  acid  and  its  salts,  called 
borates.  Of  the  latter  borax,  the  sodium  salt,  and  borocalcite 
and  colemanite,  which  are  calcium  salts,  are  the  most  im- 
portant. 

The  methods  of  preparing  boron  are  analogous  to  those  of 
making  silicon.  So  boron  may  be  prepared  by  reduction  of  its 
oxide  by  means  of  potassium,  sodium,  magnesium,  or  aluminum 
or  by  passing  the  vapors  of  boron  chloride  over  heated  sodium. 
An  amorphous  and  a  crystalline  variety  of  boron  are  known. 
The  former  results  when  the  oxide  B2O3  is  reduced  with  potas- 
sium, or  when  borax  is  heated  with  magnesium  powder. 
Amorphous  boron  is  a  brown  powder.  On  being  heated  in 
the  air  it  burns,  forming  the  oxide  B2O3  and  the  nitride  BN. 
Sulphuric  or  nitric  acid  and  other  oxidizing  agents  convert 
boron  into  boric  acid.  When  fused  with  caustic  alkalies  or 
their  carbonates,  borates  result.  Amorphous  boron  dissolves 
in  molten  aluminum,  and  on  cooling  it  crystallizes  out  in 
tetragonal  crystals,  which  are  transparent  and  generally  some- 
what colored,  due  to  impurities.  These  crystals  are  nearly  as 
hard  as  the  diamond.  They  are  less  readily  attacked  by 
reagents  than  the  amorphous  variety. 

Boron  is  trivalent  in  all  of  its  compounds.  Its  atomic  weight 
is  11.  While  boron  resembles  silicon  and  carbon  in  many 
respects,  the  formulae  of  its  compounds,  owing  to  its  trivalence, 
are  analogous  to  those  of  the  compounds  of  the  phosphorus 
group  and  to  those  of  aluminum.  The  latter  element  and 
boron  really  belong  to  the  same  family,  though  aluminum  is  a 
pronounced  metal  and  shows  but  slight  acid-forming  properties, 
while  just  the  opposite  is  true  of  boron.  The  latter  really 
occupies  a  somewhat  lone  position  amongst  the  chemical 
elements. 

Boric  Acid  and  its  Salts.  — Boric  acid  H3BO3  occurs  in  vol- 
canic regions,  particularly  in  Tuscany,  where  it  issues  from  the 
earth  in  jets  of  steam.  These  jets,  which  contain  only  small 
amounts  of  boric  acid,  are  called  soffioni,  whereas  the  hot 


SILICON  AND   BORON  301 

springs  from  which  the  jets  issue  are  termed  fumaroles.  The 
vapors  are  condensed  in  small  natural  or  artificial  basins  sur- 
rounding the  fumaroles,  and  the  boric  acid  is  finally  obtained 
by  evaporation  to  the  point  at  which  the  acid  crystallizes  out, 
the  heat  necessary  being  furnished  by  the  hot  springs.  The 
presence  of  boric  acid  in  these  steam  jets  is  due  to  the  fact 
that  boric  acid  may  be  volatilized  with  water  vapor.  In  the 
Caucasus  Mountains  and  in  some  of  the  hot  springs  of  Cali- 
fornia, boric  acid  issues  from  the  earth  in  a  similar  manner. 
Much  boric  acid  is  also  prepared  from  borax  Na2B4O7-10  H2O, 
particularly  in  Nevada  and  California.  A  hot,  concentrated 
solution  of  borax  is  treated  with  either  hydrochloric  or  sulphuric 
acid,  and  on  cooling  boric  acid  crystallizes  out.  The  reaction 
is  :  — 

Na2B407  +  5  H20  +  2  HC1  =  2  NaCl  +  4  H3BO3. 

Boric  acid  crystallizes  in  shining  white  scales  which  are 
"soapy"  to  the  touch.  At  18°,  100  parts  of  water  dissolve  3.9 
parts  of  boric  acid,  whereas  at  100°,  33  parts  of  the  acid  are 
thus  dissolved.  This  fact  makes  it  simple  to  recrystallize  boric 
acid  from  its  aqueous  solutions.  The  acid  is  quite  weak.  It 
affects  litmus  but  slightly,  and  its  taste  is  not  sour  but  simply 
astringent,  and  somewhat  bitter.  Solutions  of  boric  acid  turn 
turmeric  paper  reddish  brown.  To  bring  out  this  color  the 
paper  must  be  dried  when  very  dilute  solutions  are  used. 
This  test  for  boric  acid  is  a  very  delicate  one.  When  the  paper 
reddened  by 'boric  acid  is  treated  with  caustic  alkali,  a  black 
stain  is  produced,  which  further  serves  to  characterize  boric 
acid.  On  treating  boric  acid  with  alcohol  and  sulphuric  acid, 
a  volatile  ester,  ethyl  borate,  is  formed,  which  when  ignited 
burns  with  a  characteristic  green  flame.  This  also  serves  as  a 
test  for  boric  acid.  Boric  acid  is  often  used  in  medicine  and 
surgery  as  an  antiseptic.  It  is  also  employed  in  making  certain 
glazes  for  pottery,  and  it  is  still  sometimes  used  as  a  preserva- 
tive for  meat,  fresh  fish,  milk,  and  other  foods.  The  latter 
practice  is  to  be  condemned,  because  the  substance  is  injurious 
to  health. 

At  100°  boric  acid  loses  water  and  so  forms  metaboric  acid 
HBO2,  which  on  further  heating  to  140°  passes  over  into 
pyroboric  acid  or  tetraboric  acid  H2B4O7.  The  latter  on  igni- 


302  OUTLINES  OF   CHEMISTRY 

tion  forms  the  trioxide  or  boric  anhydride  B2O3,  which  fuses 
at  a  high  temperature  and  congeals  to  a  glassy  mass  on  cooling. 
When  treated  with  water  it  forms  boric  acid. 

Salts  of  the  acid  H3BO3  are  not  known,  but  the  esters  like 
(CH3)3BO3  and  (C2H5)3BO3  are  well  known.  Metaborates 
like  NaBO2  have  been  formed,  but  they  are  unstable. 

By  far  the  most  important  salt  of  boric  acid  is  borax 
Na2B4O7  •  10  H2O.  It  is  the  sodium  salt  of  tetraboric  acid 
H2B4O7,  which  may  be  considered  as  4H3BO3  minus  5  H2O. 
Borax  is  found  in  the  borax  lake  of  California  and  in  certain 
marshes  of  that  state  and  Nevada.  It  also  occurs  in  Thibet, 
Ceylon,  and  Bolivia.  Large  quantities  of  borax  and  boric  acid 
are  prepared  from  colemanite  Ca2B6On  •  5  H2O,  which  is  found 
in  California  and  Oregon.  The  amount  of  borax  produced 
from  the  deposits  in  the  United  States  in  1910  was  42,357 
tons. 

Borax  solutions  have  a  slightly  alkaline  reaction  toward 
indicators,  which  is  explained  by  the  fact  that  boric  acid  is 
weak  arid  its  salts  are  somewhat  decomposed  by  hydroly- 
sis. At  100°,  100  parts  of  water  dissolve  201.4  parts  of 
Na2B4O7  •  10  H2O,  whereas  at  10°  only  4.6  parts  are  thus  dis- 
solved. Borax  crystallizes  in  large  monoclinic  prisms  from 
solutions  below  50° ;  above  that  temperature  the  crystals 
formed  are  octahedra  of  the  composition  Na2B4O7 -5  H2O. 
The  salt  comes  in  the  market  in  both  forms. 

When  borax  is  heated,  it  swells  up  because  of  loss  of  water 
in  the  form  of  steam.  A  clear  liquid  is  finally  obtained  which 
solidifies  to  borax  glass.  The  latter  when  molten  dissolves 
many  metallic  oxides,  and  these  solutions  have  colors  charac- 
teristic of  the  metals  they  contain,  a  fact  that  is  often  used  in 
chemical  analysis,  and  in  making  glazes  and  enamels  for 
pottery. 

Borax  is  used  in  the  laundry  for  softening  water,  and  to 
increase  the  gloss  of  starch  in  ironing.  It  is  further  employed 
as  a  flux  in  welding  and  brazing  metals,  as  a  mordant  in  dyeing 
fabrics,  as  an  antiseptic  in  medicine,,  and  as  a  preservative.  It 
ought  not  to  be  used  as  a  preservative  for  foods. 

Other  Compounds  of  Boron.  —  Boron  hydride  BH3  is  a  gas 
formed  by  the  action  of  magnesium  boride  upon  hydrochloric 
acid. 


SILICON   AND   BORON  303 

Boron  nitride  BN  is  a  white  solid  formed  by  the  direct  union 
of  nitrogen  with  boron  when  heated.  Water  vapor  decom- 
poses it  at  high  temperatures,  forming  boric  acid  and  ammonia. 

Boron  trifluoride  BF3  is  a  colorless,  pungent  gas  made  by  the 
action  of  hydrofluoric  acid  on  boron  trioxide,  or  by  heating  the 
latter  with  fluorspar :  — 

3  CaF2  +  2  B2O3  =  Ca3B2O6  +  2  BF3. 

Boron  trichloride  BC13  is  a  colorless  liquid  of  pungent  odor. 
It  boils  at  18.2°  and  is  decomposed  by  water  into  hydrochloric 
and  boric  acids  :  — 

BC13  +  3  H2O  =  3  HC1  +  H3BO3. 

Boron  carbide  B6C  is  obtained  as  an  extremely  hard  solid  by 
heating  boron  with  carbon  in  the  electric  furnace. 

Boron  sulphide  B2S3  forms  small  white  crystals  obtained  by 
heating  boron  and  sulphur  together.  The  sulphide  is  decom- 
posed by  water  with  violence,  thus  :  — 

B2S8  +  6  H20  =  2  B(OH)3  +  3  HaS. 


CHAPTER  XIX 

PHOSPHORUS,   ARSENIC,   ANTIMONY,    AND    BISMUTH 

Occurrence  and  Preparation  of  Phosphorus.  —  Phosphorus  does 
not  occur  in  the  free  state  in  nature  because  of  its  great  affinity 
for  oxygen.  It  is  widely  distributed  in  the  form  of  phosphates, 
particularly  as  calcium  phosphate  Ca3(PO4)2,  or  apatite  3  Ca3- 
(PO4)2  -h  Ca(ClF),  though  it  is  at  times  also  found  as  wavellite 
2  A12(PO4)2  +  A12(OH)6  +  9  H2O,  vivianiteFe3(PO4)2  +  8  H2O, 
and  pyromorphite  3  Pb3(PO4)2  +  PbCl2.  In  iron  ores,  phos- 
phorus occurs  as  phosphates  of  iron  and  calcium,  and  these  are 
obtained  from  the  slags  of  blast  furnaces.  Calcium  phosphate 
is  found  in  many  rocks  and  in  all  fertile  soils.  Phosphorus  is 
also  an  essential  ingredient  of  plant  and  animal  tissues.  It 
is  specially  necessary  in  the  development  of  the  seeds  of  plants, 
hence  its  importance  in  the  soil,  from  which  the  phosphates  are 
taken  up  by  the  roots  of  plants.  The  ash  of  bones  consists  of 
80.85  per  cent  calcium  phosphate.  In  the  brain,  nerves,  blood, 
albumen,  and  muscles,  phosphorus  plays  an  important  role. 
It  occurs  here  in  complex  compounds  with  carbon,  hydrogen, 
nitrogen,  oxygen,  and  sulphur,  the  nervous  tissues  being  espe- 
cially rich  in  a  compound  called  lecithine  C42H86NPO3.  The 
urine  and  excreta  of  animals  always  contain  phosphates. 

Phosphorus  was  first  prepared  in  1669  by  the  alchemist 
Brandt,  in  Hamburg,  who  evaporated  urine  and  heated  the 
residues  mixed  with  sand  to  high  temperatures.  The  process 
was  kept  a  secret,  but  was  soon  discovered  by  Boyle  in  Eng- 
land and  Kunkel  in  Germany.  Gahn  showed  that  calcium 
phosphate  is  abundant  in  bones  (1769),  and  two  years  later 
Scheele  developed  a  method  for  preparing  phosphorus  from 
bone  ash.  Thus  calcium  sulphate  and  phosphoric  acid  are 
formed  by  means  of  the  following  reaction  :  — 

Ca3(P04)2  +  3  H2S04  =  3  CaSO4  +  2  H3PO4. 

The  calcium  sulphate  is  insoluble,  while  the  phosphoric  acid 
remains  in  solution  and  is  drained  off.     This  solution  is  evapo- 

304 


PHOSPHORUS,   ARSENIC,    ANTIMONY,   AND  BISMUTH        305 


rated  to  dryness  after  coke,  charcoal,  or  sawdust  have  been 
added,  and  the  mass  is  then  transferred  to  retorts  and  heated. 
In  this  way  water  is  driven  off  first,  finally  carbon  monoxide, 
hydrogen,  and  phosphorus  are  formed,  the  latter  being  con- 
densed and  collected  under  water.  An  older  process  consists 
of  first  forming  monocalcium  phosphate,  which  under  the  name 
of  superphosphate  is  used  as  a  fertilizer,  thus :  — 

Ca3(PO4)2  +  2  H2SO4  =  2  CaSO4  +  CaH4(PO4)2. 

This  is  then  heated  to  form  calcium  metaphosphate  Ca(PO3)2:  — 

CaH4(PO4)2  =  2  H2O  +  Ca(PO3)2. 

Finally,  by  mixing  the  calcium  metaphosphate  with  sand  and 
coke  or  charcoal,  and  heating  the  mixture  in  earthenware  re- 
torts, of  which  a  number  are  placed  in  a  furnace,  the  phospho- 
rus is  obtained  and  condensed  as  before.  The  reaction  is  :  — 

2  Ca(PO3)2  +  2  SiO2  + 10  C  =  2  CaSiO3  +  10  CO  +  4  P. 

Figure  117  shows  an  arrange- 
ment of  retorts  for  making 
phosphorus. 

By  using  the  electric  fur- 
nace, phosphorus  is  now  being 
prepared  in  a  simpler  way, 
the  process  being  a  continuous 
one.  Calcium  phosphate  is 
thoroughly  mixed  with  carbon 
and  silica  in  pulverized  form, 
and  this  mixture  is  heated  to 
a  high  temperature  in  the 
electric  furnace.  Figure  118 
shows  the  arrangement.  The 
charge  is  fed  in  continuously 
on  top  by  the  conveyor,  the  cal- 
cium silicate  slag  is  tapped  off 
at  the  bottom,  and  the  phos- 
phorus vapors  issue  from  the 
pipe  in  the  upper  part  of  the 
furnace  and  are  condensed  un- 
der water.  The  reaction  is  :  — 

Ca3(P04)2  +  3  Si02  +  50  =  3  CaSiO3  +  5  CO  +  2  P. 


1 

I 

1 

1 

1 

1 

1 

1 

1 

FIG.  117. 


306 


OUTLINES   OF  CHEMISTRY 


Thus  the  silica  lays  hold  of  the  calcium  oxide  as  it  were,  form- 
ing calcium  silicate,  and  the  oxygen  is  taken  away  from  the 

phosphorus  by  the  carbon  at 
the  high  temperature,  carbon 
monoxide  being  formed.  Phos- 
phorus when  first  condensed 
as  described  is  contaminated 
with  sand,  carbon,  and  other 
matter,  from  which  it  must  be 
freed.  This  is  accomplished 
by  melting  it  under  water  and 
straining  it,  also  under  water, 
of  course,  through  canvas 
sacks.  It  is  then  redistilled 
from  retorts  made  of  iron, 
and  cast  into  sticks  in  glass 
or  tin  molds  kept  in  cold 
water.  These  sticks  are  corn- 
ea, us.  monly  half  an  inch  in  diameter 
and  7.5  inches  long,  so  that 

nine    sticks    make    approximately    a    pound    of    phosphorus. 
Phosphorus  is  shipped  immersed  in  water  in  tin  cans. 

Properties  and  Allotropic  Forms  of  Phosphorus. — The  phos- 
phorus obtained  by  the  methods  above  described  is  known  as 
yellow  or  white  phosphorus.  It  is  a  pale  yellow,  translucent, 
waxlike  solid,  which  in  a  high  state  of  purity  is  nearly  color- 
less. In  the  cold  it  is  brittle,  somewhat  above  room  tempera- 
tures it  has  the  consistency  of  wax,  at  44°  it  melts  under  water, 
and  at  269°  it  boils  under  atmospheric  pressure.  Yellow  phos- 
phorus is  practically  insoluble  in  water,  but  it  may  be  dissolved 
to  some  extent  in  alcohol,  ether,  benzene,  and  various  ethereal 
oils  and  fats.  It  is  copiously  soluble  in  carbon  disulphide,  from 
which  it  may  be  obtained  in  rhombic  dodecahedra  of  the  iso- 
metric system  (Fig.  46)  by  evaporating  off  the  solvent  out  of 
contact  with  the  air. 

When  exposed  to  the  air,  phosphorus  slowly  oxidizes,  during 
which  process  the  oxidation  products  form  fumes,  and  emit  a 
faint  light  that  is  visible  in  the  dark.  From  the  latter  phe- 
nomenon phosphorus  derives  its  name.  By  such  slow  oxida- 
tion phosphorus  gradually  forms  a  solution  of  hypophosphoric 


PHOSPHORUS,   ARSENIC,    ANTIMONY,   AND   BISMUTH        307 

acid  which  has  reducing  properties.  During  this  oxidation  at 
room  temperatures,  ozone  and  ammonium  nitrite  are  also  formed 
from  the  air.  At  about  35°  phosphorus  catches  fire  in  the  air. 
It  must  consequently  be  kept  under  water.  If  a  little  of  the 
solution  of  phosphorus  in  carbon  disulphide  is  poured  upon 
filter  paper  and  allowed  to  evaporate,  the  finely  divided  phos- 
phorus remaining  on  the  paper  takes  fire  spontaneously. 
Phosphorus  is  very  poisonous,  0.1  gram  being  a  fatal  dose  for 
adults.  Employees  in  match  factories  are  apt  to  suffer  from 
phosphorus  poisoning,  which  manifests  itself  in  enlargement  of 
the  liver  and  necrosis  of  the  jawbones.  Phosphorus  should 
always  be  handled  with  a  forceps  and  with  great  care,  for  phos- 
phorus burns  are  dangerous  and  very  slow  to  heal. 

When  yellow  phosphorus  is  heated  from  250°  to  300°  in 
closed  vessels  out  of  contact  of  the  air,  it  is  gradually  converted 
into  red  phosphorus,  an  allotropic  form  of  phosphorus  which 
was  discovered  by  Schrotter  in  1845.  The  reaction  is  accom- 
panied with  evolution  of  heat,  and  is  never  quite  complete,  be- 
ing reversible.  When  red  phosphorus  is  heated  to  260°  in  a 
current  of  carbon  dioxide  or  nitrogen,  and  the  vapors  are  con- 
densed under  water,  the  yellow  variety  is  again  obtained. 
Light  acting  on  yellow  phosphorus  slowly  produces  some  of  the 
red  variety,  so  that  ordinary  sticks  of  phosphorus  often  have  a 
reddish  brown  outer  appearance.  Red  phosphorus  is  also  called 
amorphous;  it  does  not  emit  light  in  the  dark.  It  may  be 
heated  to  about  200°  in  the  air  without  taking  fire  and  con- 
sequently need  not  be  kept  under  water.  It  is  insoluble  in 
carbon  disulphide  and  other  solvents  that  dissolve  yellow  phos- 
phorus. Moreover,  red  phosphorus  is  not  poisonous ;  and,  in 
general,  it  is  much  less  active  than  yellow  phosphorus,  which 
contains  more  energy.  The  specific  gravity  of  red  phosphorus 
is  2.25.  By  careful  heating,  it  may  be  sublimed.  The  atomic 
weight  of  phosphorus  is  31.  Its  valence  is  either  three  or  five. 
The  vapors  of  red  and  yellow  phosphorus  are  identical.  The 
density  of  the  vapor  corresponds  to  the  formula  P4. 

Uses  of  Phosphorus,  Matches.  —  A  small  portion  of  the  phos- 
phorus produced  is  used  for  poisoning  rats  and  other  vermin. 
Most  of  it  is  used  in  making  matches.  The  annual  production 
of  phosphorus  amounts  to  over  3000  tons.  Flint,  steel,  and 
tinder  were  still  used  to  light  fires  at  the  beginning  of  the  nine- 


308  OUTLINES   OF   CHEMISTRY 

teenth  century.  In  1812  the  first  matches  made  their  appear- 
ance. They  were  invented  by  Chancel,  and  consisted  of  sticks 
dipped  in  molten  sulphur  which  was  afterwards  covered  with 
sugar  mixed  with  potassium  chlorate.  To  light  such  a  match 
its  head  was  brought  in  contact  with  concentrated  sulphuric 
acid,  which  was  commonly  absorbed  in  asbestus  and  kept  in  a 
bottle.  Thus  chloric  acid  was  liberated,  and  this  set  the  sul- 
phur and  sugar  on  fire.  In  1827  friction  or  lucifer  matches 
came  into  use.  These  had  a  head  consisting  of  potassium 
chlorate,  antimony  sulphide,  and  glue.  They  were  set  on  fire 
by  rubbing  them  vigorously  on  sandpaper.  Phosphorus 
matches  appeared  in  the  market  in  1832.  They  contained  a 
little  phosphorus  in  place  of  the  sulphide  of  antimony,  which 
caused  them  to  ignite  more  readily.  Soon  potassium  nitrate 
came  into  use  in  matches  in  place  of  potassium  chlorate,  which 
is  apt  to  cause  explosions.  At  present  the  oxidizing  agents  in 
matches  are  red  lead  Pb3O4,  lead  peroxide  PbO2,  or  manganese 
peroxide  MnO2.  In  making  matches  the  ends  of  the  well-dried 
sticks  are  first  dipped  into  paraffine.  Afterwards  they  are 
dipped  into  the  igniting  mixture,  consisting  of  phosphorus 
stirred  into  a  solution  of  glue  or  dextrine,  to  which  the  oxidiz- 
ing agents  are  added,  together  with  some  coloring  matter  like 
lamp-black,  chalk,  or  ultramarine  to  form  a  paste  of  proper  con- 
sistency. Safety  matches  were  invented  by  Bottger  in  1848. 
They  had  a  head  of  potassium  chlorate  and  antimony  trisul- 
phide  like  the  lucifer  matches,  but  it  contained  enough  glue  so 
that  the  match  ignited  with  great  difficulty  on  ordinary  sur- 
faces. However,  by  rubbing  these  matches  on  a  surface  con- 
taining red  phosphorus,  which  was  glued  on  the  box,  they 
would  ignite  very  readily.  These  safety  matches,  which  are 
often  called  Swedish  matches,  for  they  were  first  placed  011  the 
market  in  large  quantities  in  Sweden,  are  now  in  common  use. 
The  use  of  the  ordinary  match  that  will  ignite  by  friction  on 
any  surface  is  prohibited  by  law  in  some  countries.  The 
modern  safety  matches  commonly  have  a  head  consisting  of 
potassium  chlorate,  potassium  bichromate,  powdered  glass, 
and  glue  or  dextrine;  and  the  friction  surface  on  the  box 
contains  antimony  trisulphide,  red  phosphorus,  manganese 
dioxide,  and  glue.  The  purpose  of  the  powdered  glass  in 
the  head  is  to  increase  the  friction,  the  heat  from  which  raises 


PHOSPHORUS,   ARSENIC,   ANTIMONY,   AND  BISMUTH       309 

the  temperature  so  that  the  phosphorus  unites  vigorously  with 
the  oxygen  of  the  oxidizing  agents,  thus  setting  the  match 
on  fire. 

Compounds  of  Phosphorus  with  Hydrogen. — Three  compounds 
of  phosphorus  and  hydrogen  are  known.  They  are  called 
phosphines  or  phosphureted  hydrogen.  Their  composition  cor- 
responds to  the  formulae  :  — 

PH3,  a  gas  ;  P2H4,  a  liquid ;  and  P4H2,  a  solid. 

Gaseous  phosphine  PH3  is  prepared  by  heating  phosphorus  in 
a  concentrated  solution  of  caustic  potash  out  of  contact  with 
the  air.  The  reaction  is  :  — 

P4  +  3  KOH  +  3  H2O  =  3  KH2PO2  +  PH8. 

In  addition,  there  is  always  some  hydrogen  and  P2H4  formed. 
The  vapors  of  the  latter  are  spontaneously  inflammable  in  the 
air.  The  experiment  is  conducted  with  the  apparatus  shown 
in  Fig.  119.  The  small  flask  is  filled  half  full  of  caustic  potash 


FIG.  119. 

solution,  and  the  remaining  air  is  displaced  by  conducting  in 
illuminating  gas  or  hydrogen  through  the  small  tube  at  the 
left,  which  is  then  closed.  On  applying  heat,  phosphine  forms 
and  catches  fire,  forming  white  smoke  rings  as  it  issues  from 
the  mouth  of  the  delivery  tube,  which  is  kept  under  warm  water 


310  OUTLINES   OF   CHEMISTRY 


to  prevent  its  clogging  by  phosphorus 

that  mi£ht  disti11  over  and  solidify 

in  the  end  of  the  tube.  If  the  phos- 
phine  formed  is  first  passed  through 
alcohol  or  hydrochloric  acid,  the 
P2H4  is  removed  and  the  gas  PH3  is 
then  no  longer  spontaneously  inflam- 
mable in  the  air.  In  a  simpler  manner, 
phosphine  may  be  obtained  by  treating 
calcium  phosphide  with  water  or  dilute 
FIG.  120.  hydrochloric  acid  (Fig.  120)  thus  :  — 

Ca3P2  +  6  H20  =  3  Ca(OH)2  +  2  PH3,  or 
Ca3P2  +  6  HC1  =  3  CaCl2  +  2  PH3. 

In  these  reactions,  smaller  amounts  of  the  solid  and  liquid 
hydrides  of  phosphorus,  P4H2  and  P2H4,  are  also  obtained  by 
secondary  reactions.  Phosphides  of  magnesium,  zinc,  and  iron 
similarly  yield  phosphine  with  hydrochloric  acid. 

By  heating  phosphorous  or  hypophosphorous  acid,  phosphine 
is  produced,  thus  :  — 

4  H3P03      =     3  H3P04  +  PH3,  or 

phosphorous  acid  phosphoric  acid 

2  H8P02     =     H3P04    +  PH3. 

hypophosphorous  acid         phosphoric  acid 

When  phosphonium  iodide  PH4I  is  treated  with  caustic  alka- 
lies, phosphine  is  formed,  thus  :  — 

PH4I  +  NaOH  =  Nal  +  H2O  +  PH3. 
Water  also  decomposes  phosphonium  iodide  :  — 
PH4I  +  H2O  =  HI  +  H2O  +  PH3. 

Gaseous  phosphine  is  colorless.  It  boils  at  —  85°  and  solidi- 
fies at  133°.  The  gas  has  the  odor  of  rotten  fish  and  is  very 
poisonous.  Heated  to  about  100°  in  the  air  it  burns  and  forms 
water  and  phosphoric  acid.  Phosphine  is.  but  slightly  soluble 
in  water.  Alcohol  dissolves  it  more  copiously. 

With  the  hydrohalogens  phosphine  forms  phosphonium  com- 
pounds, which  are  analogous  to  ammonium  salts.  Phosphonium 
iodide,  the  best  known  of  the  phosphonium  compounds,  is  pre- 
pared by  the  following  reaction  :  — 

PH3  +  HI  =  PH4I  ;  which  is  analogous  to 


PHOSPHORUS,   ARSENIC,   ANTIMONY,   AND   BISMUTH        311 

Phosphonium  iodide  is  a  very  unstable,  colorless,  crystalline 
salt,  which  is  decomposed  by  water  into  phosphine  and  hydri- 
odic  acid,  as  stated  above.  Oxygen  acids  do  not  form  phos- 
phonium  salts  with  phosphine. 

Liquid  phosphine  P2H4  ^s  analogous  to  hydrazine  N2H4.  It 
is  a  colorless  liquid  of  specific  gravity  1.01  at  15°.  It  boils  at 
57°  and  is  insoluble  in  water. 

Solid  phosphine  P4H2  is  a  yellow,  flocculent  powder  which  is 
devoid  of  odor  and  taste.  It  does  not  dissolve  in  water.  At 
about  160°  it  takes  fire  in  the  air. 

Compounds  of  Phosphorus  with  the  Halogens.  —  Phosphorus 
forms  compounds  with  all  of  the  halogens.  These  have  the 
general  formulae  PX3  and  PX5.  The  chlorides  PC13  and  PC15 
are  the  most  important. 

Phosphorus  trichloride  PC13  is  formed  when  chlorine  is 
passed  upon  phosphorus  in  a  retort.  The  action  proceeds 
readily  with  liberation  of  heat,  the  product  being  a  colorless 
liquid  of  pungent  odor.  In  the  pure  state  phosphorus  trichlo- 
ride boils  at  76°  and  solidifies  at  —  115°.  Its  specific  gravity  is 
1.613  at  0°.  Water  decomposes  it  :  — 

PC13  +  3  H20  =  3  HC1  +  P(OH)3. 

Phosphorus  pentachloride  PC15  is  formed  by  treating  phos- 
phorus trichloride  with  chlorine,  or  by  passing  an  excess  of 
chlorine  upon  phosphorus  in  a  retort  :  — 

PC18+C12  =  PC16,  or 


=  4PC15. 

The  product  is  a  light  yellow,  finely  crystalline  solid  which  can- 
not be  melted  under  atmospheric  pressure,  for  the  temperature 
at  which  its  vapor  tension  equals  atmospheric  pressure  lies  be- 
low, the  melting  point  of  the  compound.  Under  the  pressure 
of  its  own  vapor  in  a  sealed  tube,  phosphorus  pentachloride 
may  be  melted  at  148°.  When  heated,  phosphorus  pentachlo- 
ride decomposes  into  phosphorus  trichloride  and  chlorine  :  — 


At  300°  this  dissociation  is  nearly  complete.  The  action  is 
reversible,  as  indicated.  It  is  quite  similar  to  the  dissociation 
of  ammonium  chloride  by  means  of  heat:  — 


312  OUTLINES   OF  CHEMISTRY 

With  water,  phosphorus  pentachloride  forms  hydrochloric 
acid  and  phosphorus  oxychloride  :  — 

PC15  +  H2O  =  2  HC1  +  POC18. 

The  latter  is  a  colorless  liquid  of  specific  gravity  1.712  at  0°. 
It  boils  at  107.5°  and  melts  at  1.8°.  On  further  treatment  with 
water,  the  oxychloride  also  decomposes,  yielding  hydrochloric 
acid  and  phosphoric  acid  :  — 

POC13  +  3  H2O  =  3  HC1  +  H3PO4. 

Phosphorus  trifluoride  PF3  is  a  colorless  gas.  It  boils  at 
-  95°  and  congeals  at  -  160°.  The  pentafluoride  PF6  melts 
at  —  83°  and  boils  at  —  75°.  These  compounds  are  decom- 
posed by  water  like  the  analogous  chlorides,  but  more  slowly. 
Phosphorus  oxyfluoride  POF3  is  a  gas  which  may  be  liquefied  at 
-50°. 

Phosphorus  tribromide  PBr8  is  a  colorless  liquid  boiling  at 
172°.  Its  specific  gravity  is  2.925  at  0°.  Phosphorus  penta- 
bromide  PBr6  forms  yellow  crystals,  which  on  heating  disso- 
ciate into  bromine  and  phosphorus  tribromide. 

Phosphorus  triodide  PI3  forms  dark  red,  prismatic  crystals 
melting  at  61°.  Phosphorus  pentaiodide  is  not  known,  but  a 
diphosphorus  tetraiodide  P2T4  is  known.  It  forms  orange-yellow 
crystals  which  melt  at  110°. 

On  treatment  with  water,  both  the  bromides  and  iodides  of 
phosphorus  are  decomposed  into  the  hydrohalogen  acids  and 
oxygen  acids  of  phosphorus.  From  phosphorus  pentabro- 
mide,  phosphorus  oxybromide  POBr3  may  be  obtained  in  a 
manner  analogous  to  the  formation  of  POC13.  The  treatment 
of  phosphorus  tribromide  or  triodide  with  water  affords  excellent 
methods  for  making  pure  hydrobromic  and  hydriodio  acids,  as 
already  stated. 

Oxides  and  Acids  of  Phosphorus. — The  following  oxides  of 
phosphorus  are  well  known  :  phosphorus  trioxide  P2O3 ; 
phosphorus  tetroxide  P2O4;  and  phosphorus  pentoxide  P2O5. 
Of  these  the  latter  is  the  most  important  by  far.  The  trioxide 
is  a  white  crystalline  solid  melting  at  22.5°.  It  is  obtained 
together  with  t^e  pentoxide  by  burning  phosphorus  in  an  in- 
sufficient amount  of  oxygen.  The  tetroxide  is  a  white  solid 
formed,  together  with  red  phosphorus,  by  heating  the  trioxide 


PHOSPHORUS,   ARSENIC,    ANTIMONY,   AND  BISMUTH       313 

in  a  sealed  tube  to  440°.  Phosphorus  pentoxide  P2O5  *s  formed 
when  phosphorus  is  burned  in  the  air  or  in  oxygen.  It  is  a 
light  white  powder  which  unites  with  water  with  great  avidity, 
forming  metaphosphoric  acid,  thus :  — 

P205  +  H20  =  2HP03. 

Phosphorus  pentoxide  is  the  best  drying  agent  known.  Its 
action  on  water  is  accompanied  with  evolution  of  much  heat 
and  a  hissing  noise  resembling  that  accompanying  the  quench- 
ing of  hot  iron. 

In  union  with  different  amounts  of  water,  phosphorus  pen- 
toxide forms  three  acids,  thus  .  — 

P2O5  +  H2O  =  2  HPO3  (metaphosphoric  acid), 
P2O5  -f-  2  H2O  =  H4P2O7  (pyrophosphoric  acid), 
P2Q5  +  3  H2O  =  2  H3PO4  (orthophosphoric  acid). 

By  union  with  two  or  six  molecules  of  water  phosphorus  tri- 
oxide  forms  two  acids,  thus  :  — 

2  P2O3  +  2  H2O  =  4  HPO2  (metaphosphorous  acid), 
2  P2O3  +  6  H2O  =  4  H3PO3  (phosphorous  acid). 

There  are  also  known  hypophosphoric  acid  H4P2O6  and  hypo- 
phosphorous  acid  H3PO2.  The  former  is  prepared  by  allowing 
sticks  of  phosphorus  to  oxidize  slowly  in  contact  with  moist 
air,  under  which  conditions  phosphoric  and  phosphorous  acids 
are  also  formed  to  some  extent.  The  acid  is  tetrabasic,  and 
consequently  is  able  to  form  four  kinds  of  salts  by  successive 
replacement  of  the  hydrogen  atoms.  Hypophosphorous  acid 
H3PO2  may  be  liberated  from  its  barium  salt  by  action  of  sul- 
phuric acid,  thus  :  — 

8  P  +  3  Ba(OH)2  +  6  H2O  =  2  PH3  +  3  Ba(H2PO2)2,  and 
Ba(H2P02)2  +  H2S04=  BaSO4  +  2  H3PO2. 

It  is  a  monobasic  acid,  forming  crystals  that  melt  at  17.4°.  It 
is  a  strong  reducing  agent.  On  being  heated,  it  yields  phos- 
phine  and  phosphoric  acid. 

Orthophosphoric  Acid.  —  This  compound  is  also  called  simply 
phosphoric  acid.  Its  composition  is  expressed  by  the  formula 
H3PO4.  It  may  be  considered  as  derived  from  the  hypothet- 
ical pentahydroxide  of  phosphorus  P(OH)6  by  loss  of  one 
molecule  of  water.  Pure  phosphoric  acid  is  prepared  by  action 


314  OUTLINES   OF   CHEMISTRY 

of  phosphorus  pentoxide  on  water  or  by  the  oxidation  of  phos- 
phorus by  means  of  nitric  acid.  Phosphoric  acid  is  also  made 
by  the  action  of  sulphuric  acid  upon  calcium  phosphate.  The 
calcium  sulphate  formed  simultaneously,  being  insoluble,  is 
readily  removed  and  the  clear  solution  containing  the  phos- 
phoric acid  is  then  evaporated.  It  commonly  still  contains 
some  calcium  salts  which  may  be  precipitated  by  means  of 
alcohol.  Solutions  of  pure  phosphoric  acid  may  be  evaporated 
to  a  thick,  colorless  sirup  of  specific  gravity  1.88,  from  which 
upon  cooling  a  crystalline  mass  is  obtained,  which  melts  at  42°. 
The  crystals  are  deliquescent  and  dissolve  in  water  with  great 
readiness.  The  solutions  are  strongly  acidic  in  character. 
The  acid  is  not  poisonous.  Phosphoric  acid  is  tribasic  and  con- 
sequently is  able  to  form  three  classes  of  salts,  the  primary,  sec- 
ondary, and  tertiary  phosphates,  for  instance  :  — 

H3P04  +  KOH  =  KH2P04  +  H2O. 
H3P04  +  2  KOH  =  K2HP04  +  2  H2O. 
H3P04  +  3  KOH  =  K3P04  +  3  H2O. 

The  tertiary  phosphates  are  the  normal  or  neutral  salts  ;  whereas 
the  secondary  and  primary  salts  still  contain  one  and  two  hydro- 
gen atoms  respectively  in  the  molecule.  The  hydrogen  atoms 
need  not  all  be  replaced  by  the  same  metal  or  radical.  Thus 
we  have  sodium  ammonium  hydrogen  phosphate  NaNH4HPO4, 
which  is  also  known  as  microcosmic  salt.  Magnesium  ammonium 
phosphate  MgNH4PO4  forms  white  insoluble  crystals  ;  it  is  of 
importance  in  analytical  chemistry.  Solutions  of  the  secondary 
salts  have  an  alkaline  reaction,  being  to  some  extent  decom- 
posed by  hydrolysis.  The  tertiary  salts  are  much  more  hy- 
drolyzed  by  water;  indeed  they  are  stable  only  as  solids,  and 
are  obtained  by  evaporating  the  acid  to  dryness  with  the  proper 
amount  of  alkali.  These  salts  are  not  decomposed  by  heat, 
whereas  both  the  secondary  and  primary  phosphates  lose  water 
on  being  heated  ;  so,  for  instance  :  — 

2Na2HPO4:£Na4P207  +  H20;  and 

sodium 
pyrophosphate 

NaH2P04^±NaP03  +  H2O. 

sodium 
metaphosphate 


PHOSPHORUS,    ARSENIC,    ANTIMONY,    AND  BISMUTH        315 

Thus  secondary  phosphates  yield  pyrophosphates,  and  primary 
phosphates  yield  metaphosphates,  on  heating.  Conversely,  on 
treatment  with  water  the  pyrophosphates  gradually  pass  back 
into  secondary  phosphates,  and  the  metaphosphates  into  pri- 
mary phosphates.  Microcosmic  salt  and  magnesium  ammonium 
phosphate  lose  ammonia  as  well  as  water  on  being  heated, 
thus:  — 

NaNH4HP04  =  H2O  +  NH3  +  NaPO3. 

2  MgNH4P04  =  H20  +  2  NH3  +  Mg2P2O7. 

Pyrophosphoric  acid  H4P2O7  is  formed  by  heating  phosphoric 
acid  to  about  250°  till  a  sample  neutralized  with  ammonia  and 
tested  with  silver  nitrate  solution  yields  a  white  precipitate. 
The  white  precipitate  is  silver  pyrophosphate  Ag4P2O7,  whereas 
the  phosphate  of  silver  Ag3PO4  is  yellow.  The  formation  of 
pyrophosphoric  acid  takes  place  thus  :  — 

2  H3P04  =  H4P207  +  H20. 

The  aqueous  solutions  of  pyrophosphoric  acid  are  fairly  stable, 
the  acid  passing  over  into  orthophosphoric  acid  but  slowly. 
The  presence  of  sulphuric  or  nitric  acids  hastens  the  change. 
Though  the  molecule  of  pyrophosphoric  acid  contains  four 
hydrogen  atoms,  but  two  kinds  of  pyrophosphates  are  known. 
These  correspond  to  the  types  K4P2O7  and  K2H2P2O7. 

By  the  color  of  the  silver  salt,  pyrophosphoric  acid  is  readily 
distinguished  from  orthophosphoric  acid.  From  metaphos- 
phoric  acid,  pyrophosphoric  acid  is  distinguished  by  the  fact 
that  it  does  not  coagulate  albumen  like  the  former. 

Metaphosphoric  acid  HPO3  is  made  by  heating  phosphoric  acid 
to  400°: 

H3P04=H20  +  HPO3; 

or  by  treating  phosphorus  pentoxide  with  water; 


or  by  heating  ammonium  phosphate; 

(NH4)2HP04  =  2  NH3  +  H20  +  HPO3. 

The  acid  is  a  glassy,  semitransparent  mass  which  is  also  called 
glacial  phosphoric  acid.  In  contact  with  water,  it  slowly  passes 
over  into  phosphoric  acid,  the  action  being  hastened  by  boiling. 
The  acid  is  monobasic  and  is  analogous  to  nitric,  chloric,  and 


316 


OUTLINES   OF   CHEMISTRY 


bromic  acids.  Solutions  of  glacial  phosphoric  acid  coagulate 
albumen  and  give  white  precipitates  with  the  chlorides  of 
barium  or  calcium,  which  behavior  is  different  from  that  of 
solutions  of  pyrophosphoric  acid. 

Phosphorous  acid  H3PO3  forms  as  one  of  the  products  of  the 
slow  oxidation  of  phosphorus  in  moist  air.  It  is  best  prepared 
by  treating  phosphorus  trichloride  with  water  and  driving  off 
the  hydrochloric  acid  formed  simultaneously,  by  heating  to 
180°.  The  acid  forms  very  hygroscopic  crystals  that  melt  at 
70°.  On  heating,  it  decomposes  into  phosphoric  acid  and 
phosphine :  — 

4H8P03=3H3P04  +  PH3. 

At  the  high  temperature  at  which  the  reaction  takes  place,  the 
phosphoric  acid  formed  passes  over  into  metaphosphoric  acid, 
and  the  phosphine  burns  with  a  green  flame.  Though  phos- 
phorous acid  has  three  hydrogen  atoms  in  the  molecule,  it  is 
only  dibasic.  Its  salts  correspond  to  the  type  Na2HPO3,  the 
third  hydrogen  atom  not  being  replaceable  by  a  metal. 

Formulae  of  the  Acids  of  Phosphorus.  —  The  following  struc- 
tural formulae  of  the  oxy-acids  of  phosphorus  will  serve  to 
impress  their  relationships  further  :  — 


— H  orthophosphoric 
Q_H  acid. 

\0-H 


O-H 


\ 


,       ,      . 
pyrophosphoric 


metaphosphoric 


\)_H          acid' 


O 


p>/H          phosphorous 
\\0-H          acid. 
\O-H 


/O-H 

P-O-H 
Nvri 


^  hypophosphoric 

p//O-H  acid. 


\ 


hypophosphorous 
acid. 


O-H 


PHOSPHORUS,   ARSENIC,    ANTIMONY,   AND   BISMUTH        317 

It  will  be  seen  that  the  dibasic  character  of  phosphorous  acid 
is  expressed  by  connecting  the  non-replaceable  hydrogen  atom 
directly  with  the  phosphorus.  Similarly  the  monobasic  char- 
acter of  hypophosphorous  acid  is  indicated  by  connecting  the 
two  non-replaceable  hydrogen  atoms  directly  with  phosphorus. 

Compounds  of  Phosphorus  with  Sulphur.  —  With  sulphur, 
phosphorus  unites  directly,  forming  a  series  of  compounds  : 
P4S3,  P2S3,  P3S6,  and  P2S5.  The  action  of  yellow  phosphorus 
upon  hot  sulphur  is  violent;  the  sulphides  are  consequently 
made  by  using  red  phosphorus.  Phosphorus  pentasulphide 
P2S5  forms  yellow  crystals  which  melt  at  275°.  The  liquid 
boils  at  518°.  With  potassium  sulphide  it  forms  potassium 
sulphophosphate  :  — 


With    phosphorus     pentachloride    phosphorus     sulphochloride 
PSC1    results:  — 


The  latter  compound  is  a  colorless  liquid  of  specific  gravity 
1.168  at  0°.  It  boils  at  125°,  and  decomposes  upon  treatment 
with  water  :  — 

PSC18  +  4  H2O  =  3  HC1  +  H2S  +  H3PO4. 

Occurrence,  Preparation,  and  Properties  of  Arsenic.  —  Arsenic 
is  very  widely  distributed  in  nature  in  minute  quantities.  It 
rarely  occurs  in  the  uncombined  state,  being  found  in  larger 
quantities  in  combination  with  sulphur,  as  in  realgar  As2S2  and 
orpiment  As2S3.  It  is  also  found  combined  with  oxygen,  as  in 
arsenolite  As2O3,  and  with  iron  and  sulphur  and  cobalt  and 
sulphur,  as  in  arsenical  pyrites  or  mispickel  FeAsS  and 
cobaltite  CoAsS. 

Arsenic  is  commonly  prepared  by  heating  mispickel  or  by 
reducing  arsenolite  with  carbon.  The  reactions  are  :  — 

FeAsS  =  FeS  +  As. 
2  As2O3  +  6  0  =  As4  +  6  CO. 

Arsenic  is  volatile  ;  it  sublimes,  and  is  readily  condensed. 

Arsenic  is  steel-gray  in  color,  has  a  bright  metallic  luster, 
and  is  very  brittle.  Its  specific  gravity  is  5.73  at  15°.  On 
heating,  it  volatilizes  without  melting  ;  but  under  pressure  it 
may  be  melted  at  about  480°.  At  450°  its  vapor  tension  equals 


318  OUTLINES   OF   CHEMISTRY 

atmospheric  pressure.  Heated  in  the  air,  it  burns,  the  fumes 
having  a  garlic-like  odor  and  the  flame  a  pale  lavender  color  ; 
these  are  characteristic  of  arsenic.  Between  560°  and  860°  the 
vapor  of  arsenic  is  about  150  times  as  heavy  as  hydrogen.  Hence 
the  molecular  weight  is  approximately  300;  and  since  the  atomic 
weight  of  arsenic  is  75  as  determined  from  the  analysis  of  the 
chloride,  the  molecular  formula  of  arsenic  is  As4.  Between  1600° 
and  1700°,  Victor  Meyer  found  the  vapor  of  arsenic  to  be  only 
75  times  as  heavy  as  hydrogen,  which  leads  to  the  molecular 
formula  As2.  The  valence  of  arsenic  is  either  three  or  five, 
and  the  formulae  of  its  compounds  are  consequently  analogous 
to  those  of  nitrogen  and  phosphorus. 

Arsenic  burns  to  As2O3  in  the  air  when  heated  to  180°.  It 
combines  directly  with  many  elements  like  chlorine,  bromine, 
sulphur,  and  some  of  the  metals.  When  boiled  with  nitric 
acid  or  aqua  regia,  arsenic  is  oxidized  to  arsenic  acid  H3AsO4. 

Besides  the  metallic  form  of  arsenic  above  described,  this 
element  may  be  obtained  as  yellow  crystals  by  rapidly  cooling 
its  vapor.  The  crystals  resemble  ordinary  phosphorus  in  that 
they  dissolve  in  carbon  bisulphide.  Arsenic  itself  does  not  act 
as  a  poison,  for  it  is  not  taken  up  by  the  animal  system.  Its 
insoluble  sulphides  also  are  not  especially  toxic  in  character. 
However,  all  other  compounds  of  arsenic,  notably  arsine  AsH3, 
arsenious  oxide  As2O3,  halogen  compounds,  and  salts  of  arsenious 
and  arsenic  acids  are  very  poisonous.  From  0.1  to  0.4  gram  of 
arsenious  oxide  is  sufficient  to  cause  death.  The  antidote  for 
arsenic  is  freshly  precipitated  ferric  hydroxide. 

Arsine,  Arseniureted  Hydrogen,  AsH3.  —  Arsine  is  a  colorless 
gas.  It  was  discovered  by  Scheele  in  1755.  It  melts  at 
—  113.5°  and  boils  at  —  55°.  It  is  analogous  to  ammonia  NH3 
and  phosphine  PH3.  It  is  commonly  prepared  (1)  by  the 
action  of  hydrochloric  or  sulphuric  acid  upon  the  arsenide  of  zinc 
or  sodium,  or  (2)  by  introducing  compounds  of  arsenic  in  a 
flask  containing  zinc  and  hydrochloric  or  sulphuric  acid.  The 
reactions  involved  in  these  processes  are  typified  by  the  follow- 
ing equations  :  — 

(1)  Zn8  As2  +  6  HC1  =  3  ZnCl2  +  2  AsH3. 
AsNa8  +  3  H2SO4  =  3  NaHSO4  +  AsH8. 

(2)  As2O3  +  12  H  =  3  H2O  +  2  AsH3. 


PHOSPHORUS,    ARSENIC,   ANTIMONY,    AND   BISMUTH       319 

The  odor  of  arsine  is  very  disagreeable,  resembling  that  of 
garlic.  Arsine  is  extremely  poisonous  and  great  care  must  con- 
sequently be  exercised  in  experimenting  with  it.  Arsine  does  not 
unite  with  water  or  with  acids ;  it  thus  exhibits  much  less  basic 
properties  than  ammonia  or  phosphine.  Ignited  in  the  air, 
arsine  burns  with  a  pale  lavender  flame,  forming  water  and 
arsenious  oxide  :  — 

2  AsH3  +  3  O2  =  3  H2O  +  As2O3. 

On  being  heated,  the  gas  readily  dissociates  into  arsenic  and 

hydrogen  :  — 

4AsH3  =  As4-f  6H2. 

So  when  dry  arsine  is  passed  through  a  tube  heated  to  dull 
redness,  the  reaction  just  given  takes  place,  the  arsenic  con- 
densing in  form  of  a  metallic  mirror  in  the  colder  parts  of  the 
tube.  Since  solutions  of  all  arsenic  compounds  when  intro- 
duced into  a  flask  containing  zinc  and  hydrochloric  or  sulphuric 
acid  yield  arsine,  a  simple  and  very  efficient  method  of  testing 
arsenic,  known  as  Marsh's  test,  has  been  devised.  The  appara- 
tus is  shown  in  Fig.  121.  Pure  zinc  and  hydrochloric  acid  are 


FIG.  121. 

introduced  into  the  flask.  The  calcium  chloride  in  the  tube 
serves  to  dry  the  gases  evolved.  After  all  air  has  been  ex- 
pelled, the  hydrogen  is  lighted  and  the  solution  to  be  tested 
for  arsenic  is  poured  down  the  funnel  tube.  If  arsenic  is 
present,  the  flame  will  acquire  the  characteristic  pale  lavender 
color,  and  dark  spots  of  metallic  arsenic  will  be  deposited  upon 
a  white  porcelain  dish  held  in  the  flame.  If  the  tube,  which 
should  be  of  hard  glass,  is  heated  as  shown,  a  mirror  of  metallic 


320  OUTLINES   OF   CHEMISTRY 

arsenic  will  deposit  on  the  sides  of  the  tube  just  beyond  the 
flame.  Both  the  mirror  and  the  spots  are  soluble  in  sodium 
hypochlorite  or  bleaching  powder  solution.  It  is  to  be  noted 
that  compounds  of  antimony  under  like  treatment  yield  similar 
spots  and  mirrors  ;  these  are,  however,  not  soluble  in  hypo- 
chlorites.  Moreover,  the  arsenical  mirror  is  more  volatile  than 
that  of  antimony.  The  former  may  be  converted  into  yellow 
sulphide  of  arsenic  and  the  latter  into  red  sulphide  of  antimony 
by  means  of  hydrogen  sulphide. 

When  conducted  into  a  solution  of  silver  nitrate,  arsine  pre- 
cipitates metallic  silver,  thus  :  — 

2  AsH3  +  12  AgNO3  +  3  H2O  =  As2O3  +  12  HNO3  +  12  Ag. 

Since  the  corresponding  antimony  hydride,  stibine  SbH3,  does 
not  reduce  silver  nitrate  solutions  thus,  this  reaction  may  be 
used  to  distinguish  between  arsine  and  stibine. 

Compounds  of  Arsenic  with  the  Halogens.  —  Of  these  com- 
pounds arsenic  trichloride  AsCl3  is  the  most  important.  There 
are  also  known  :  the  trifluoride  AsF3,  boiling  at  63°  and  melting 
at  —  8.5°;  the  tribromide  AsBr3,  melting  at  31°  and  boiling  at 
221° ;  the  triodide  AsT3,  melting  at  140°,  as  well  as  iodides  of 
the  formulae  AsI2  and  AsI5. 

Arsenic  trichloride  is  formed  by  conducting  chlorine  upon 
powdered  arsenic  contained  in  a  retort,  or  by  the  action  of 
hydrochloric  acid  upon  arsenic  trioxide.  It  is  a  colorless,  fum- 
ing liquid  of  specific  gravity  2.205  at  0°.  It  boils  at  129°  and 
solidifies  to  a  crystalline  mass  at  — 18°.  It  is  very  poisonous. 
Water  decomposes  it :  — 

2  AsCl3  +  3  H20  =  As203  +  6  HC1. 

By  addition  of  concentrated  hydrochloric  acid,  the  hydrolysis 
may  be  reversed.  It  will  be  recalled  that  this  cannot  be  done 
in  the  case  of  the  analogous  chloride  of  phosphorus. 

Oxides  and  Oxy- acids  of  Arsenic.  —  Two  oxides,  the  trioxide 
As2O3  and  the  pentoxide  As2O5,  are  known  ;  and  the  corre- 
sponding acids,  arsenious  acid  H3AsO3  and  arsenic  acid  H3AsO4, 
are  of  importance. 

Arsenic  trioxide  As2O3,  also  called  "white  arsenic"  or  com- 
monly simply  "  arsenic"  is  the  commonest,  and  by  far  the  most 
important,  of  all  the  compounds  of  arsenic.  It  is  found  in  nature 


PHOSPHORUS,    ARSENIC,    ANTIMONY,    AND   BISMUTH        321 

and  is  formed  when  arsenic  burns  in  the  air  or  in  oxygen. 
Arsenic  trioxide  is  manufactured  on  a  commercial  scale  by 
roasting  arsenical  pyrites  in  the  air.  In  this  process  iron  oxide 
remains  as  a  non-volatile  residue,  sulphur  dioxide  escapes,  and 
the  arsenious  oxide  condenses  as  a  white  powder  upon  the  brick 
waljs  of  the  chambers.  It  is  purified  by  resublimation.  In 
the  year  1910,  1497  tons  of  arsenious  oxide  were  produced  in 
the  United  States.  About  four  times  this  amount  is  annually 
produced  in  Europe.  On  heating  arsenic  trioxide,  it  gradually 
forms  an  amorphous  glassy  mass,  which  after  a  time  becomes 
white,  crystalline,  and  opaque.  Below  200°  the  crystals  formed 
are  octahedra  of  the  regular  system,  whereas  above  that  tem- 
perature crystallization  in  monoclinic  forms  takes  place.  At 
800°  the  vapor  density  of  arsenious  oxide  corresponds  to  the 
formula  (As2O3)2,  whereas  at  about  1800°  the  density  of  the 
gas  leads  to  the  simple  formula  As2O3,  the  double  molecules 
having  been  dissociated.  Arsenic  trioxide  is  readily  reduced 
to  arsenic  by  heating  it  with  carbon,  or  cyanide  of  potassium. 
Its  conversion  to  arsine  has  already  been  mentioned.  In  water 
it  dissolves  but  slightly.  Hydrochloric  acid  dissolves  it,  form- 
ing arsenic  trichloride. 

The  trioxide  has  a  sweetish,  disagreeable  taste.  It  is  a  strong 
poison.  It  is  used  as  rat  poison,  also  in  taxidermy,  in  calico 
printing,  in  the  manufacture  of  certain  kinds  of  glass,  in  the 
preparation  of  many  other  compounds  of  arsenic,  and  in  medi- 
cine. Freshly  precipitated  ferric  hydroxide  forms  an  insoluble 
compound  with  arsenious  oxide  and  is  consequently  used  as  an 
antidote  in  cases  of  poisoning. 

Arsenious  acid  H3AsO3  has  not  been  isolated.  It  probably 
exists  in  the  aqueous  solutions  of  arsenious  oxide.  Its  salts, 
the  arsenites,  are  known.  Among  these  may  here  be  mentioned 
silver  arsenite  Ag3AsO3  and  copper  hydrogen  arsenite,  or 
Scheele's  green,  CuHAsO3.  Salts  of  meta-arsenious  acid  HAsO2 
are  also  known,  like  KAsO2  and  Pb(AsO2)2.  Paris  green,  also 
called  Schweinfurt  green,  is  a  double  salt  of  cupric  arsenite  and 
cupric  acetate  Cu3As2O6-Cu(C2H3O2)2.  It  is  used  as  a  poison 
for  potato  bugs  and  other  insects. 

Arsenic  acid  H3AsO4  is  readily  produced  by  oxidation 
of  arsenious  acid.  Scheele  prepared  arsenic  acid  in  1775 
by  passing  chlorine  into  arsenic  trioxide  suspended  in 


322  OUTLINES  OF  CHEMISTRY 

water ;  nitric  acid  or  a  mixture  of  nitric  and  hydrochloric 
acids  serves  equally  well.  The  reaction  in  the  former 
case  is :  — 

As203  +  2  C12  +  5  H20  =  2  H3  AsO4  +  4  HC1. 

The  acid  forms  rhombic,  deliquescent  prisms  or  plates  of  the 
composition  2  H3AsO4  +  H2O.  At  100°  the  water  of  crystalli- 
zation passes  off.  At  about  180°  the  acid  loses  water,  passing 
over  into  pyroarsenic  acid  H4As2O7,  which  on  being  heated  still 
further  again  loses  water,  forming  meta-arsenic  acid  HAsO3. 
So  far  then  the  behavior  is  entirely  similar  to  that  of  phos- 
phoric acid,  though  in  contact  with  water  pyro-  and  meta- 
arsenic  acids  at  once  form  arsenic  acid.  On  further  ignition  of 
meta-arsenic  acid,  water  is  again  split  off  and  arsenic  pentoxide 
As2O5  is  formed,  thus  :  — 

2  HAs03  =  H20  +  As205. 

It  will  be  recalled  that  metaphosphoric  acid  cannot  thus  be 
decomposed  into  P2O5  and  water.  Furthermore,  phosphorus 
pentoxide  is  very  stable  when  heated,  whereas  arsenic  pentoxide 
decomposes  upon  ignition  into  arsenic  trioxide  and  oxygen  :  — 

As2O5  =  As2O3  +  O2. 

The  salts  of  arsenic  acid  are  quite  analogous  to  those  of 
phosphoric  acid.  Thus,  there  are  primary,  secondary,  and  tertiary 
arsenates,  also  pyroarsenates  and  meta-ar senates.  In  contact 
with  water,  however,  all  the  salts  form  orthoarsenates  at  once. 

Sulphides  of  Arsenic.  —  Three  sulphides  of  arsenic  are  known, 
namely :  the  disulphide  As2S2,  the  trisulphide  As2S3,  and  the 
pentasulphide  As2S5. 

Arsenic  disulphide  As2S2  occurs  in  nature  as  realgar,  in  red 
monoclinic  prisms.  It  is  also  manufactured  by  fusing  sulphur 
and  arsenic  together.  Thus  made,  it  forms  a  dark  red,  glassy 
substance,  which  in  pulverized  condition  is  sometimes  used  as  a 
pigment  in  paints.  A  mixture  of  1  part  arsenic  disulphide, 
12  parts  saltpeter,  and  3.5  parts  sulphur  when  ignited  makes 
white  Bengal  fire. 

Arsenic  trisulphide  As2S3  occurs  in  nature  in  short  rhombic 
prisms"  as  orpiment.  It  was  formerly  used  as  a  pigment.  It 
is  readily  obtained  as  a  lemon-yellow  precipitate  by  passing 


PHOSPHORUS,   ARSENIC,   ANTIMONY,   AND   BISMUTH       323 

hydrogen  sulphide  into  a  solution  of  arsenic  trioxide  in  hydro- 
chloric acid  :  — 

2  AsCl3  +  3  H2S  =  As2S3  +  6  HC1. 

On  heating  the  precipitate  with  concentrated  hydrochloric  acid, 
it  may  be  redissolved ;  that  is,  the  reaction  just  given  may  be 
reversed.  Arsenic  trisulphide  may  also  be  obtained  by  fusing 
sulphur  and  arsenic  together  in  the  right  proportions.  In  am- 
monium sulphide,  arsenic  trisulphide  is  soluble,  forming  ammo- 
nium sulpharsenite  (NH4)3  AsS3 :  — 

As2S3  +  3  (NH4)2S  =  2  (NH4)3AsS3. 

In  solution  of  yellow  ammonium  sulphide,  that  is,  in  ammonium 
sulphide  containing  an  excess  of  sulphur,  arsenic  trisulphide 
dissolves  as  ammonium  sulpharsenate  (NH4)3AsS4 :  — 

As2S3  +  3  (NH4)2S  +  28  =  2  (NH4)3AsS4. 

On  treatment  with  hydrochloric  acid  the  sulpharsenites' and  sulph- 
arsenates  are  decomposed :  — 

2  (NH4)3AsS3  +  6  HC1  =  As2S3  +  6  NH4C1  +  3  H2S. 
2  (NH4)3  AsS4  +  6  HC1  =  As2S5  +  6  NH4C1  +  3  H2S. 

Arsenic  pentasulphide  As2S5,  made  by  means  of  the  reaction 
just  given  or  by  melting  together  sulphur  and  arsenic  in  proper 
proportions,  is  a  yellow  solid  which  may  be  sublimed  when 
heated  out  of  contact  with  the  air. 

Occurrence,  Preparation,  and  Properties  of  Antimony.  —  Anti- 
mony (stibium)  is  sometimes,  though  rarely,  found  in  nature 
in  the  uncombined  state.  When  .thus  found,  it  occurs  in 
rhombohedral  crystals.  The  mineral  stibnite  Sb2S3,  found  in 
Hungary  and  Japan,  is  the  chief  source  of  antimony,  though 
the  latter  also  occurs  combined  with  sulphur  in  many  native 
sulphides  of  lead,  copper,  silver,  iron,  and  arsenic.  Native 
oxide  of  antimony,  senarmontite  Sb2O3,  forming  white  octa- 
hedra  of  the  regular  system,  is  also  known.  Stibnite  was 
known  in  ancient  times.  The  Chaldeans  manufactured  vari- 
ous articles  out  of  metallic  antimony,  and  the  alchemists 
frequently  used  the  metal. 

Antimony  is  prepared  by  heating  stibnite  with  iron,  thus :  — 


324  OUTLINES   OF   CHEMISTRY 

It  is  also  made  by  roasting  stibnite  in  the  air,  and  reducing 
the  tetroxide  thus  formed,  by  means  of  carbon.  The  reactions 
are  as  follows  :  — 

Sb2S3  +  5  02  =  3  S02  +  Sb204. 
Sb2O4  +  4  C  =  4  CO  +  2  Sb. 

To  free  the  antimony  "  regulus "  so  obtained  from  iron,  lead, 
copper,  etc.,  it  is  fused  with  a  little  sulphur  or  saltpeter. 
Thus  the  impurities  are  converted  to  sulphides  or  oxides,  which 
float  on  top  and  can  be  removed.  Antimony  free  from  arsenic 
and  other  metals  may  be  obtained  by  reducing  pure  sodium 
metantimoniate  NaSbO3. 

Antimony  is  a  hard,  brittle,  silvery-white  metal  having  a 
high  metallic  luster.  It  can  readily  be  ground  to  powder. 
At  625°  it  melts,  and  on  cooling  it  forms  rhombohedral  crystals. 
Its  boiling  point  is  approximately  1400°,  and  its  specific  gravity 
is  6.75.  In  the  air  it  remains  practically  unchanged,  but  when 
strongly  heated  it  burns  with  a  bluish  white  flame  to  Sb2O3  or 
Sb2O4.  Introduced  into  an  atmosphere  of  chlorine,  it  takes  fire 
and  burns  to  SbCl5.  It  dissolves  in  hot  concentrated  sulphuric 
acid,  also  in  aqua  regia,  but  nitric  acid  converts  it  into  Sb2O3 
or  antimonic  acid  H3SbO4.  Hydrochloric  acid  acts  slowly  on 
antimony,  liberating  hydrogen.  The  latter  gas  is  also  formed 
by  the  action  of  steam  on  antimony  at  high  temperatures. 

The  atomic  weight  of  antimony  is  120.2.  The  vapor  density 
leads  to  a  molecular  weight  of  approximately  290,  which  rep- 
resents a  formula  lying  between  Sb2  and  Sb3.  The  valence  of 
antimony  is  either  three  or  five.  Its  compounds  consequently 
have  formulae  analogous  to  those  of  nitrogen,  phosphorus,  and 
arsenic.  The  latter  is  a  close  relative  of  antimony. 

Metallic  antimony  is  much  used  in  alloys,  particularly  in  type 
metal  and  britannia  metal.  Type  metal  consists  of  approxi- 
mately 25  per  cent  antimony,  25  per  cent  tin,  and  50  per  cent 
lead.  The  presence  of  antimony  in  alloys  makes  them  hard. 
Furthermore,  antimony  expands  as  it  congeals  (resembling 
water  in  this  behavior)  and  consequently  fills  molds  perfectly, 
thus  yielding  sharply  defined  castings. 

Hydrogen  Antimonide,  Stibine,  SbH3.  —  This  compound  is 
analogous  to  ammonia,  phosphine,  and  arsine.  It  is  quite  simi- 
lar to  the  latter  and  is  prepared  by  similar  methods.  So,  for 


PHOSPHORUS,    ARSENIC,   ANTIMONY,    AND   BISMUTH        325 

instance,  by  treating  an  alloy  of  magnesium  and  antimony  or 
zinc  and  antimony  with  dilute  hydrochloric  or  sulphuric  acid, 
stibine  is  formed.  Again,  by  introducing  a  solution  of  any 
antimony  compound  into  a  flask  in  which  zinc  is  being  acted 
upon  by  hydrochloric  or  sulphuric  acid,  stibine  results,  which 
in  this  case  is  mixed  with  hydrogen. 

Stibine  is  a  colorless  gas  of  peculiar  odor,  reminding  one 
somewhat  of  that  of  hydrogen  sulphide.  The  odor  is  distinctly 
different  from  that  of  arsine.  Stibine  melts  at  —  88°  and  boils 
at  — 17°.  The  gas  readily  dissociates  into  antimony  and  hydro- 
gen, thus :  — 

2  SbH3  =  2  Sb  +  3  H2. 

The  change  begins  at  150°.  Even  when  diluted  with  hydro- 
gen, stibine  is  largely  decomposed  when  passed  through  a  tube 
heated  to  150°,  yielding  a  deposit  of  antimony  in  the  form  of 
a  mirror,  which  is  insoluble  in  hypochlorites.  Thus,  in  the 
apparatus  used  for  making  Marsh's  test  for  arsenic,  antimony 
compounds  would  yield  a  similar  mirror ;  but  the  latter  is 
readily  distinguished  from  arsenic  by  the  method  described 
under  arsine.  The  dissociation  of  stibine  is  practically  com- 
plete at  200°,  at  which  temperature  arsine  remains  unchanged. 
Stibine  is  moderately  poisonous.  Water  dissolves  about  four 
times  its  own  volume  of  the  gas  at  room  temperature.  In  the 
air  or  in  oxygen,  when  ignited,  stibine  burns  with  a  bluish 
white  flame,  forming  water  and  Sb2O3.  Conducted  into  a 
silver  nitrate  solution,  stibine  is  decomposed,  the  antimony 
being  precipitated  as  silver  antimonide  SbAg3.  When  pure 
or  when  diluted  with  hydrogen,  stibine  may  be  kept  unchanged  ; 
but  the  presence  of  even  small  amounts"  of  oxygen  in  the  gas 
leads  to  the  deposition  of  some  of  the  antimony. 

Compounds  of  Antimony  with  the  Halogens.  —  Of  these,  anti- 
mony trichloride  SbCl3  and  antimony  pentachloride  SbCl5  are 
of  most  importance. 

Antimony  trichloride  is  formed  by  the  action  of  chlorine  on 
antimony  or  of  hydrochloric  acid  on  antimony  sulphide  :  — 

2  Sb  +  3  C12  =  2  SbClg. 
Sb2S3  +  6  HC1  =  3  H2S  +  2  SbCl3. 

The  antimony  trichloride  is  purified  by  distillation.  It  is  a 
colorless  crystalline  mass  which  at  ordinary  temperatures  is 


326  OUTLINES   OF   CHEMISTRY 

soft,  reminding  one  of  the  consistency  of  butter,  hence  it  goes 
by  the  name  of  butter  of  antimony.  It  melts  at  73°  and  boils 
at  223°.  At  26°  its  specific  gravity  is  3.064.  Its  vapor  is 
229  times  as  heavy  as  hydrogen,  which  fact  leads  to  the 
formula  SbCl3.  It  is  deliquescent  and  has  caustic  properties. 
Antimony  trichloride  is  used  as  a  mordant,  also  in  medicine 
and  in  burnishing  metals,  notably  gun  barrels,  to  which  it 
imparts  a  brown  hue.  Antimony  trichloride  may  be  dissolved 
in  water  containing  hydrochloric  acid.  But  when  treated  with 
water  alone,  antimony  trichloride  is  decomposed  into  hydro- 
chloric acid  and  insoluble  oxychlorides,  the  composition  of 
which  varies  according  to  the  temperature  and  relative  amount 
of  water  used.  Two  oxychlorides  of  antimony,  SbOCl  and 
(SbOCl)2  •  Sb2O3,  are  well  known  as  white  crystalline  powders. 
They  are  formed  thus  :  — 

(1)  SbCl3  +  H20  =  SbOCl  +  2  HC1. 

(2)  4  SbCl3  +  5  H2O  =  (SbOCl)2  •  Sb2O3  +  10  HC1. 

The  second  reaction  takes  place  in  hot  solutions.  The  com- 
pound (SbOCl)2»  Sb2O3,  or  Sb4O5Cl2,  was  used  by  the  Italian 
physician  Victor  Algarotus,  and  is  consequently  known  as  the 
powder  of  algaroth. 

Antimony  pentachloride  is  prepared  by  burning  antimony  in 
an  excess  of  chlorine  or  by  conducting  chlorine  upon  antimony 
trichloride.  It  is  a  fuming  liquid  of  yellow  color.  At  —  6° 
its  crystals  melt.  It  can  only  be  distilled  in  a  partial  vac- 
uum, for  on  heating  it  readily  dissociates  into  chlorine  and 
the  trichloride.  With  water  it  forms  crystalline  hydrates, 
SbCl5-H2O  and  SbOl5-4H2O.  Antimony  pentachloride  is 
decomposed  by  hot  water.  It  readily  gives  off  part  of  its 
chlorine,  and  is  consequently  used  in  organic  chemistry  in 
chlorinating  substances.  It  will  be  observed  that  while 
antimony  pentachloride  forms  crystalline  hydrates  with 
water,  the  latter  decomposes  the  chlorides  of  phosphorus  at 
once. 

Antimony  trifluoride  SbF3  forms  deliquescent  rhombic  crystals 
that  are  not  decomposed  by  cold  water.  With  ammonium  sul- 
phate it  forms  a  compound  that  is  used  as  a  mordant. 

Antimony  pentafluoride  SbF6  is  an  amorphous  gummy  mass. 
It  readily  enters  into  the  formation  of  double  salts. 


PHOSPHORUS,   ARSENIC,    ANTIMONY,   AND   BISMUTH       327 

Antimony  tribromide  SbBr3  forms  white  rhombic  crystals  that 
melt  at  94°.  The  salt  boils  at  275°,  and  is  decomposed  by 
water. 

Antimony  triiodide  SbI3  forms  three  different  varieties  of 
crystals.  The  common  red  crystals  melt  at  171°.  The  boiling 
point  is  430°. 

Antimony  pentiodide  SbI5  is  a  dark  brown,  crystalline  mass 
of  melting  point  79°.  It  is  unstable. 

Oxides  and  Oxy-acids  of  Antimony.  —  There  are  three  oxides 
of  antimony :  antimony  trioxide  Sb2O3,  antimony  tetroxide 
Sb2O4,  and  antimony  pentoxide  Sb2O5.  The  trioxide  acts 
mainly  as  a  base,  though  toward  very  strong  bases,  like  caustic 
potash  and  soda,  it  is  also  able  to  act  as  an  acid.  The  tetroxide 
exhibits  neither  acid  nor  basic  properties,  whereas  the  pentoxide 
acts  solely  in  an  acid-forming  capacity. 

Antimony  trioxide  is  found  in  nature  as  senarmontite.  It  is 
formed  by  burning  antimony  in  the  air  or  by  oxidizing  the 
metal  with  nitric  acid.  The  oxide  is  white  and  may  be  sub- 
limed. It  crystallizes  in  octahedra  or  rhombic  prisms,  being 
dimorphous.  At  1560°  the  density  of  its  vapor  corresponds  to 
the  formula  Sb4O6,  nevertheless  it  is  commonly  called  the  tri- 
oxide. It  is  possible  that  at  higher  temperatures  it  would 
dissociate  into  Sb2O3  like  the  corresponding  oxide  of  arsenic. 
In  water  and  nitric  or  sulphuric  acid,  antimony  -trioxide  is 
practically  insoluble,  while  in  hydrochloric  or  tartaric  acid,  or 
in  acid  potassium  tartrate  or  caustic  alkalies,  it  dissolves, 
thus : — 

Sb203  4-  6  HC1  =  2  SbCl3  +  3  H2O. 
Sb2O3  +  2  KOH  =  2  KSbO2  +  H2O. 
Sb2O3  +  2  (C4H406)HK  =  2  (C4H4O6)SbO  •  K  +  H2O. 

The  salt  KSbO2  is  potassium  metantimonite.  It  is  plainly  a 
salt  of  metantimonious  acid  HSbO2,  which  may  be  considered 
as  derived  from  antimonious  acid  H3SbO3  by  loss  of  a  molecule 
of  water.  The  salt  (C4H4O6)  •  SbO  •  K  is  potassium  antimonyl 
tartrate  or  tartar  emetic.  It  contains  the  univalent  antimonyl 
group,  —  Sb  =  O,  which  is  frequently  found  in  other  antimony 
salts.  Tartar  emetic  has  been  known  for  a  long  time.  The 
salt  crystallizes  with  half  a  molecule  of  crystal  water,  a  part  of 
which  escapes  on  exposure  to  the  air.  The  salt  is  still  some- 


328  OUTLINES   OF   CHEMISTRY 

times  used  in  medicine.  Antimony  salts  were  formerly  fre- 
quently prescribed  by  physicians.  These  compounds  gained 
in  prominence  through  the  work  of  Basil  Valentine,  who  in 
the  fifteenth  century,  published  his  book  on  "  The  Triumphal 
Chariot  of  Antimonium."  The  compounds  (C4H4O6)  •  AsO  •  K 
potassium  arsenyl  tartrate  and  (C4H4O6)  •  BO  •  K  potassium 
boryl  tartrate  are  analogous  to  tartar  emetic. 

On  treating  tartar  emetic  with  dilute  sulphuric  acid,  the 
hydrate  H3SbO3  separates  out  as  a  precipitate,  which,  how- 
ever, loses  water  and  forms  metantimonious  acid  HSbO2,  i.e. 
SbO-  OH. 

The  basic  properties  of  antimony  are  shown  in  its  salts,  in 
which  either  Sb(OH)3  or  SbO  •  OH  act  as  bases.  Thus,  there 
are  known  antimony  nitrate  Sb(NO3)3,  antimony  sulphate 
Sb2(SO4)3,  and  the  halogen  salts  like  SbCl3;  farther,  when 
these  salts  are  acted  upon  by  water,  oxy-salts  or  basic  salts  are 
produced,  which  may  be  considered  as  derived  from  SbO  •  OH. 
So  antimonyl  nitrate  SbO  •  NO3  and  antimonyl  sulphate 
(SbO)2SO4  are  known,  and  antimony  oxychloride  and  tartar 
emetic,  already  mentioned,  belong  in  this  category. 

Antimony  tetroxide  is  a  white  powder  obtained  by  burning 
antimony  in  oxygen  or  by  heating  the  trioxide  in  the  air.  In 
water  it  is  insoluble,  'while  boiled  with  cream  of  tartar  it  is 
converted  into  tartar  emetic  and  metantimonic  acid,  thus  :  — 

Sb204  +  (C4H406)HK  =  (C4H4O6)SbO  -  K  +  HSbO3. 
The  tetroxide  is  also  obtained  on  igniting  antimony  pentoxide  :  — 


Antimony  tetroxide  may  be  regarded  as  the  antimonyl  salt  of 
metantimonic  acid,  which  is  HSbO3.  The  antimonyl  salt  would 
have  the  formula  (SbO)  •  SbO3. 

Antimonic  acid  H3SbO4  is  formed  as  an  insoluble  white 
powder  by  the  action  of  concentrated  nitric  acid  upon  antimony, 
or  by  the  action  of  water  on  antimony  pentachloride.  Salts  of 
this  acid  and  also  of  its  dehydration  products,  pyro-  and  metan- 
timonic acids  are  known.  So  on  fusing  antimony  with  potas- 
sium nitrate,  there  is  formed  with  explosive  violence  potassium 
metantimonate  KSbO8,  which  on  being  heated  with  water  passes 
into  solution  as  potassium  antimonate  KH2SbO4.  On  fusing 


PHOSPHORUS,   ARSENIC,    ANTIMONY,    AND   BISMUTH        329 

potassium  metantimonate  with  caustic  potash,  the  pyroantimo- 
nate  K4Sb2O7  results  :  — 

2  KSb03  +  2  KOH  =  K4Sb207  +  H2O. 
Potassium  pyroantimonate  is  decomposed  by  water :  — 
K4Sb207  +  2  H20  =  2  KOH  +  KaH2SbaO7. 

When  the  latter  salt  is  added  to  a  solution  of  a  sodium  salt, 
sodium  pyroantimonate  Na2H2Sb2O7  is  precipitated.  This  is 
practically  the  only  sodium  salt  known  that  does  not  dissolve  in 
water  readily. 

Antimonic  acid  and  its  dehydration  products  are  then  quite 
analogous  to  those  of  the  corresponding  phosphorus  and  arsenic 
compounds. 

Antimony  pentoxide  Sb2O5  is  a  yellow  powder  obtained  by 
heating  antimonic  acid  to  275°'  At  higher  temperatures  it  is 
decomposed,  yielding  the  tetroxide  and  oxygen.  With  strong 
bases  it  forms  salts.  It  is  soluble  in  hydrochloric  acid. 

Compounds  of  Antimony  with  Sulphur.  —  It  has  already  been 
mentioned  that  antimony  trisulphide  Sb2S3  is  found  in  nature 
as  stibnite.  Precipitated  from  solutions  of  antimony  salts  by 
means  of  hydrogen  sulphide,  antimony  trisulphide  is  an  orange- 
red  powder,  which  is  insoluble  in  dilute  hydrochloric  acid,  but 
soluble  in  concentrated  hydrochloric  acid,  with  concomitant 
evolution  of  hydrogen  sulphide.  In  ammonium  sulphide  it 
dissolves,  yielding  ammonium  sulphantimonite,  thus :  — 

Sb2S3  +  3  (NH4)2S  =  2  (NH4)3SbS3. 
The  latter  is  decomposed  by  hydrochloric  acid  :  — 

2  (NH4)3SbS3  +  6  HC1  =  6  NH4C1  +  Sb2S3  +  3  H2S. 

In  yellow  ammonium  sulphide,  antimony  trisulphide  dissolves 
more  readily,  yielding  ammonium  sulphantimonate  :  — 

Sb2S3  +  3  (NH4)2S  +  S2  =  2  (NH4)3SbS4. 

On  treating  the  latter  with  hydrochloric  acid,  antimony  penta- 
sulphide  Sb2S5  is  obtained  :  — 

2  (NH4)3SbS4  +  6  HC1  =  6  NH4C1  +  Sb2S6  +  3  H2S. 

Antimony  pentasulphide  may  also  be  obtained  by  treating  anti- 
monic acid  with  hydrogen  sulphide,  thus :  — 

2  H3Sb04  +  5  H2S  =  Sb2S6  +  8  H2O. 


330  OUTLINES   OF   CHEMISTRY 

It  is  a  powder  of  golden  yellow  color,  hence  it  is  called  sulphur 
auratum.  On  being  heated,  it  gives  off  sulphur  and  forms  the 
trisulphide.  In  soluble  sulphides  of  the  metals  it  dissolves, 
forming  sulphantimonates.  Thus  with  sodium  sulphide  it 
forms  Na3SbS4  +  9  H2O,  which  is  known  as  "  Schlippe's  salt." 

Antimony  pentasulphide  is  used  in  making  red  vulcanized 
caoutchouc.  The  trisulphide  is  used  in  making  matches,  also 
as  a  pigment.  Antimony-  cinnabar,  kermes  mineral,  a  mixture 
of  the  trisulphide  and  trioxide  of  antimony,  is  used  in  medicine. 

Occurrence,  Preparation,  and  Properties  of  Bismuth.  —  This 
element,  though  not  abundant  or  widely  distributed  in  nature, 
has  been  known  since  the  fifteenth  century,  when  it  was  referred 
to  by  Basil  Valentine,  who,  on  account  of  its  brittleness,  re- 
garded it  as  a  half  metal.  Bismuth  generally  occurs  in  the  free 
state  in  nature,  and  is  almost  always  fairly  pure.  Sometimes  it 
is  found  as  the  sulphide,  bismuth  glance  Bi2S3,  more  rarely  as 
the  oxide,  bismuth  ocher  Bi2O3.  The  sulphide  is  roasted  to 
oxide,  which  is  then  reduced  with  charcoal.  The  bismuth  so 
obtained,  or  the  native  bismuth,  is  refined  by  fusing  it  with 
saltpeter  or  soda  plus  a  little  potassium  chlorate.  Thus,  arsenic 
and  other  impurities,  consisting  mainly  of  lead,  iron,  antimony, 
copper,  sulphur,  etc.,  are  oxidized  and  removed  as  a  slag  that 
floats  on  the  surface.  ' 

Bismuth  is  a  white,  brittle  metal  having  a  high  metallic 
luster  and  a  slightly  reddish  sheen,  which  readily  distinguishes 
it  from  antimony.  Bismuth  is  crystalline.  Its  crystals  belong 
to  the  rhombohedral  division  of  the  hexagonal  system.  Its 
specific  gravity  is  9.82.  It  melts  at  269°,  and  may  be  distilled 
in  a  vacuum  at  about  995°.  It  is  a  rather  poor  conductor  of 
heat  and  electricity,  as  compared  with  other  metals.  The 
atomic  weight  of  bismuth  is  208,  and  its  valence  is  commonly 
either  three  or  five ;  so  that  the  formulae  of  its  compounds  are 
analogous  to  those  of  nitrogen,  phosphorus,  arsenic,  and  anti- 
mony. Nevertheless,  bismuth  is  more  pronouncedly  basic  in 
character  than  these,  and  consequently  it  is  to  be  grouped  with 
the  metals. 

In  the  air  bismuth  remains  practically  unchanged.  On 
ignition  in  the  air  it  burns  with  a  bluish- white  flame  ;  the  prod- 
uct formed  is  a  yellow  powder,  the  trioxide,  Bi2O8.  In  nitric 
acid,  bismuth  may  readily  be  dissolved,  forming  the  nitrate 


PHOSPHORUS,    ARSENIC,    ANTIMONY,    AND  BISMUTH        33i 

Bi(NO3)3 ;  likewise  when  the  metal  is  treated  with  sulphuric 
acid,  the  sulphate,  Bi2(SO4\  is  formed.  Hydrochloric  acid 
scarcely  attacks  bismuth.  The  latter  does  not  combine  with 
hydrogen. 

Bismuth  is  used  in  pharmaceutical  preparations.  It  is  also 
used  in  making  alloys  that  have  a  low  melting  point.  Of  these 
the  following  are  frequently  used :  Rose's  metal,  consisting  of 
1  part  tin,  1  part  lead,  and  2  parts  bismuth,  melts  at  93.8°; 
Newton's  metal,  consisting  of  3  parts  tin,  5  parts  lead,  and  8 
parts  bismuth,  melts  at  94.5°;  and  Wood's  metal,  which  consists 
of  1  part  tin,  2  parts  lead,  1  part  cadmium,  and  4  parts  bismuth, 
melts  at  60.5°.  On  changing  from  the  liquid  to  the  solid  state 
bismuth  expands  even  more  than  antimony.  It  is  consequently 
also  employed,  like  the  latter,  in  alloys  for  stereotyping  and 
other  purposes  where  castings  of  sharp  outline  are  required. 

Halogen  Compounds  of  Bismuth.  —  In  these  compounds  bis- 
muth is  always  trivalent.  Bismuth  chloride  BiCl3  is  made  by 
the  action  of  chlorine  upon  bismuth,  or  by  dissolving  the  latter 
in  nitro-hydrochloric  acid.  It  may  also  be  obtained  by  dissolv- 
ing the  trioxide,  Bi2O3,  in  hydrochloric  acid.  The  salt  con- 
sists of  white  crystals  melting  at  227°,  and  boiling  at  about 
445°.  It  is  soluble  in  hydrochloric  acid  solutions,  from  which 
it  is  precipitated  in  the  form  of  bismuth  oxychloride  BiOCl :  — 

BiCl3  +  H20  =  BiOCl  +  2  HC1. 

Bismuth  fluoride  BiF3  is  a  grayish  powder  formed  by  the 
action  of  hydrofluoric  acid  on  bismuth  trioxide.  On  treatment 
with  much  water,  bismuth  oxyfluoride  BiOF  is  formed.  Bis- 
muth bromide  BiBr3  forms  orange-colored  crystals  melting  at 
215°  and  boiling  at  453°.  With  water  they  yield  bismuth 
oxybromide  BiOBr.  Bismuth  iodide  BiI3  consists  of  dark 
brown  or  black  crystals  of  metallic  luster,  melting  at  439°. 
On  boiling  with  water  they  are  decomposed,  yielding  red  crys- 
tals of  bismuth  oxyiodide  BiOI. 

Halogen  compounds  of  bismuth  in  which  the  element  has  a 
valence  of  rive  have  not  been  prepared,  but  a  dichloride  of  the 
formula  (BiCl2)2  has  been  described  as  a  white  powder  formed 
by  heating  bismuth  with  mercurous  chloride. 

Oxides  of  Bismuth.  —  Bismuth  trioxide  Bi2O3,  which  is  formed 
as  a  yellow  powder  when  the  metal  is  burned  in  the  air,  is  the 


332  OUTLINES   OF  CHEMISTRY 

most  important  of  the  oxides.  It  acts  only  as  a  base,  forming 
salts  which  may  be  considered  as  derived  from  either  Bi(OH)3 
or  BiO  •  OH. 

Bismuth  dioxide  Bi2O2  is  obtained  as  a  dark  brown  precipi- 
tate by  pouring  a  solution  containing  stannous  chloride  and 
bismuth  chloride  into  caustic  potash  solution. 

Bismuth  tetroxide  Bi2O4  is  a  reddish  yellow  powder  formed 
by  heating  the  pentoxide  to  about  165°. 

Bismuth  pentoxide  Bi2O5  is  an  unstable  brown  powder  ob- 
tained by  passing  chlorine  into  caustic  potash  solution  contain- 
ing bismuth  trioxide  in  suspension.  On  being  heated,  it  forms 
the  tetroxide.  With  hydrochloric,  acid  it  forms  bismuth  tri- 
chloride and  chlorine :  — 

Bi2O5  + 10  HC1  =  5  H2O  +  2  BiCl3  +  2  C12. 

Bismuth  Salts  of  Oxy-acids.  —  The  salts  of  bismuth  with  the 
halogens  have  already  been  described.  With  sulphuric  acid 
bismuth  forms  bismuth  sulphate  Bi2(SO4)3,  which  on  treat- 
ment with  water  yields  the  oxysulphate  or  bismuthyl  sulphate 
(BiO)2SO4,  thus:- 

Bi2(SO4)3  +  4  H2°  =  (BiO)2SO4  4-  2  H2SO4  +  2  H2O. 
With  nitric  acid,  bismuth  forms  the  nitrate  Bi(NO3)3,  which 
crystallizes  in  triclinic  forms  with  rive  molecules  of  water. 
The  salt  is  decomposed  into  basic  nitrates  by  treatment  with 
water.  The  composition  of  these  basic  nitrates  varies  with  the 
temperature  and  the  relative  amounts  of  water  and  normal 
nitrate  used  in  preparing  them.  Thus  a  white  powder,  bismuth 
oxynitrate  BiO  •  NO3,  is  known.  On  boiling  this  salt  with  water, 
a  more  basic  salt  of  approximately  the  composition  BiO  •  NO3  -f 
BiO  •  OH  is  obtained  which  is  used  as  a  cosmetic  and  antiseptic 
under  the  name  bismuth  subnitrate.  Furthermore,  it  is  very 
often  prescribed  in  medicine  in  cases  of  dysentery  and  other 
disturbances  of  the  digestive  tract.  In  the  treatment  of  dis- 
eases of  the  skin,  particularly  in  cases  of  acute  inflammations, 
it  is  also  frequently  employed. 

All  salts  of  bismuth  may  be  regarded  as  derived  from  the  two 
basic  hydroxides  Bi(OH)3  and  BiO  •  OH.  The  univalent  radi- 
cal —  Bi  =  O,  bismuthyl,  is  analogous  to  the  antimonyl  radical 
—  Sb  =  O.  The  tendency  to  form  oxy-salts  or  basic  salts  is 
very  characteristic  of  bismuth  and  also  of  antimony. 


PHOSPHORUS,    ARSENIC,   ANTIMONY,   AND  BISMUTH       333 


Bismuth  Trisulphide  Bi2S3  occurs  in  nature  as  bismuth 
glance.  It  may  also  be  obtained  as  a  very  dark  brown  or  black 
precipitate  by  passing  hydrogen  sulphide  into  a  solution  of  a  salt 
of  bismuth :  — 

2  BiCl3  +  3  H2S  =  6  HC1  +  Bi2S3. 

It  is  insoluble  in  ammonium  sulphide  solution,  also  in  solutions 
of  the  sulphides  of  the  alkalies.  This  behavior  distinguishes  it 
from  the  sulphides  of  arsenic  and  antimony,  which  readily  dis- 
solve in  alkali  sulphides  as  sulpho-salts.  On  heating  a  precipi- 
tate of  bismuth  trisulphide  suspended  in  a  solution  of  an  alkali 
sulphide  to  200°,  the  compound  becomes  crystalline.  Bismuth 
trisulphide  may  also  be  obtained  by  melting  together  sulphur 
and  bismuth  in  proper  proportions. 

A  compound  of  the  composition  Bi2S2,  bismuth  disulphide, 
has  also  been  described  as  consisting  of  steel-gray  needles 
formed  by  melting  sulphur  and  bismuth  together  in  the  pro- 
portions represented  by  the  formula. 

General  Considerations  of  the  Group.  Nitrogen,  phosphorus, 
arsenic,  antimon}^,  and  bismuth  form  another  natural  group  of 
elements.  Their  atomic  weights  increase  in  the  order  named, 
and  their  physical  properties  show  a  corresponding  gradation 
of  changes,  as  is  evident  from  the  following  table :  — 


ELEMENT 

ATOMIC 
WEIGHT 

COLOR 

SPECIFIC  GRAVITY 

MELTING 
POINT 

BOILING 
POINT 

Nitrogen,  N 

14.01 

colorless 

0.885  (liquid) 

-210.5° 

-194.4° 

Phosphorus,  P 

31.0 

yellow  or 

1.8-2.3 

+  44.4° 

+  278.0° 

red 

Arsenic,  As 

75.0 

gray, 

5.7 

500° 

450° 

lustrous 

(approx.) 

(approx.) 

Antimony,  Sb 

120.2 

white, 

6.8 

625° 

1500° 

lustrous 

(approx.) 

Bismuth,  Bi 

208.0 

reddish 

9.8 

268° 

1600° 

white 

(approx.) 

The  chemical  properties  of  the  members  of  the  group  also 
present  an  interesting  series  of  changes  as  the  atomic  weight 
increases.  The  compounds  with  hydrogen  have  the  formula 
RH8.  So  we  have  ammonia  NH3,  phosphine  PH3,  arsine  AsH3, 


334 


OUTLINES   OF   CHEMISTRY 


and  stibine  SbH3.  The  stability  of  these  compounds  dimin- 
ishes in  the  order  named,  a  hydride  of  bismuth  being  unknown. 
Ammonia  has  strong  basic  properties ;  these  are  also  exhibited 
by  phosphine,  but  to  a  lesser  degree,  in  the  phosphonium  salts. 
But  arsine  and  stibine  are  no  longer  able  to  unite  with  acids  to 
form  salts.  Hydrazine  (NH2)2  has  its  analogue  in  liquid  phos- 
phine (PH2)2,  while  analogous  hydrides  of  arsenic  and  anti- 
mony are  unknown.  Furthermore,  hydrazoic  acid  HN3  and 
solid  phosphine  P4H2  stand  alone,  no  analogous  compounds  of 
the  group  being  known. 

In  general,  as  the  atomic  weight  of  the  elements  of  this  group 
increases,  the  affinity  for  hydrogen  decreases. 

Just  the  reverse  is  true  of  the  affinity  of  nitrogen,  phosphorus, 
arsenic,  antimony,  and  bismuth  for  the  halogens.  The  halogen 

III  V 

compounds  have  the  general  types  RX3  and  RX5.  Thus,  we 
have  the  following  series  of  the  halogen  compounds :  — 

HALOGEN  COMPOUNDS  OF   THE   NITROGEN   GROUP 


NF3(?) 



NC13 

_____ 

NBr3 

_____ 

_____ 

NI3+NH3 



PF3 

PF5 

PC13 

PC15 

PBr8 

PBr5 

P2I4 

PI3 



AsF3 



AsCl3 



AsBr3 



As2I4 

AsI3 



SbFs 

SbF5 

SbCl8 

SbCl5 

SbBr3 





SbI3 

Sbl, 

BiF3 



BiCl3 



BiBr3 





BiI3 



While  the  halogen  compounds  of  nitrogen  are  so  unstable  as 
to  be  explosive  in  character,  the  phosphorus  halides  possess  a 
considerable  degree  of  stability,  which  increases  as  we  pass  to 
corresponding  compounds  of  arsenic,  antimony,  and  bismuth  in 
the  order  named.  The  phosphorus  halides  are  at  once  decom- 
posed by  water  completely.  The  arsenic  halides  suffer  such 
hydrolysis  more  slowly,  and  even  incompletely  if  but  little 
water  is  used,  while  the  halides  of  antimony  and  bismuth  are 
but  partially  decomposed  by  water,  forming  oxy-salts  that  are 
fairly  stable.  These  oxy-salts  generally  have  the  formula 

ROX,  like  SbOCl,  etc.,  though  on  treatment  with  boiling 
water  they  form  more  basic  salts  because  of  further  hydrolysis. 
The  affinity  of  the  elements  of  this  group  toward  oxygen 
and  sulphur  also  diminishes  as  the  atomic  weight  increases. 
With  oxygen  we  have  the  following  compounds :  — 


PHOSPHORUS,   ARSENIC,    ANTIMONY,   AND   BISMUTH       335 


N2O 

NO 

NA 

(N02) 

^A 





PA 

(P02)2 

PA 





As2O3 



As205 





Sb2O3 

(Sb02)2 

Sb2O5 



(BiO)2 

Bi203 

(Bi02)2 

BiA 

The  oxy-acids  are  as  follows,   those   in   parentheses  being 
known  only  in  the  form  of  salts :  — 


H2N202 

HNO2 

(HAsO2) 
HSbO2 

H3P02 

H3PO3 
H3As63 

H3Sb03 

H3P04 
H3AsO4 
H3SbO4 

H4P2Or 
H4As2Or 

H4Sb2O7 

HN03 
HP03 
HAsO3 
HSbO3 
(HBi03) 



The  sulphides  are  commonly  of  the  general  type 

V 

R2S5.     They  are  given  in  the  following  table  :  — 


or 


<NAV 



NA 

As2S2 

As2S3 

As2S5 



Sb2S3 

Sb2S5 

Bi2S2 

Bi2S3 



The  sulphides  of  arsenic  and  antimony  unite  with  ammonium 
sulphide  and  other  alkaline  sulphides  to  form  sulpho-salts, 
while  the  sulphide  of  bismuth  does  not  do  this. 

Bismuth  is  already  a  fairly  pronounced  metal.  It  is,  after 
all,  not  very  closely  related  to  the  other  members  of  the  group. 
"Nitrogen,  too,  it  must  be  admitted,  stands  rather  remotely 
related  to  the  other  members  of  the  group.  Its  relation  to 
the  latter  is  more  like  that  of  fluorine  toward  the  other  halo- 
gens, or  like  oxygen  toward  sulphur,  selenium,  and  tellurium. 

Vanadium,  Columbium,  and  Tantalum.  —  These  are  rare  metal- 
lic elements  which  form  compounds  that  are  analogous  to  those 
of  the  phosphorus  group,  though  in  some  respects  these  three 
elements  are  also  similar  to  aluminum,  iron,  chromium,  and 
tungsten. 

Vanadium  (V  —  51.0)  was  discovered  by  Del  Rio  in  1801 
in  vanadinite  PbCl2  •  3  Pb3(VO4)4,  but  was  characterized  defi- 


336  OUTLINES  OF   CHEMISTRY 

nitely  as  an  element  in  1830  by  Sefstrom.  In  1867  Roscoe 
prepared  vanadium  by  heating  its  diehloride,  VC12,  in  a  current 
of  hydrogen.  Thus  made,  vanadium  is  a  crystalline,  grayish 
powder  of  specific  gravity  5.8,  which  readily  burns  to  VaO5 
in  oxygen.  Other  oxides  are  V2O,  V2O2,  V2O3,  V2O4.  The 
chlorides  are  VCla,  VC13,  VOC13,  VC14.  The  vanadates  are 
salts  of  vanadic  acid  H3VO4.  As  in  the  case  of  the  phos- 
phates, there  are  ortho-,  meta-,  and  pyrovanadates  ;  of  these 
the  metavanadates,  like  NaVO3,  are  met  most  frequently. 

Columbium  (Cb  —  93.5)  is  also  called  niobium,  Nb.  It 
occurs  together  with  tantalum  (Ta  — 181.0)  in  columbites 
and  tantalites.  These  metals  may  be  prepared  like  vanadium. 
Moissan  prepared  tantalum  by  reduction  of  its  oxide  Ta2O5 
with  carbon  in  the  electric  furnace.  Tantalum  is  a  steel-gray, 
malleable  metal  which  melts  at  about  2300°.  It  is  now  being 
used  to  make  filaments  for  incandescent  electric  lamps.  As 
the  metal  conducts  better  than  carbon,  tantalum  filaments 
must  be  made  longer  than  carbon  filaments  to  obtain  the 
necessary  amount  of  electrical  resistance. 


CHAPTER   XX 

CLASSIFICATION  OF  THE  ELEMENTS 
THE  PERIODIC  SYSTEM 

WE  have  already  seen  that  the  chemical  elements  may  be 
divided  into  metals  and  non-metals,  although  a  sharp  line  of 
division  between  these  two  groups  does  not  exist.  It  would 
be  natural  to  classify  the  elements  according  to  their  physical 
and  chemical  properties.  Experience  has  shown  that  the 
properties  of  the  elements  are  closely  related  to  their  atomic 
weights.  As  early  as  1817  Dobereiner  called  attention  to  the 
fact  that  the  atomic  weight  of  strontium,  87.62,  is  approxi- 
mately the  arithmetical  mean  of  the  atomic  weights  of  barium, 
137.37,  and  calcium,  40.09.  These  three  metals  are  very  sim- 
ilar in  character,  forming  the  group  of  the  alkaline  earth 
metals.  A  number  of  other  elements  that  are  closely  related 
also  form  similar  groups  of  three,  or  so-called  triads,  which  is 
evident  from  the  following  cases  :  — 

Chlorine,  35.46 ;  Bromine,  79.92  ;  Iodine,  126.92. 

35.46  +  126.92 

T  31.19. 

Sulphur,  32.07;  Selenium,  79.2;  Tellurium,  127.5. 

32.07  +  127.5  _  7Q  _ 

~2~ 

Phosphorus,  31.0;  Arsenic,  75.0;  Antimony,  120.2. 

31.0  +  120.2      7.  „ 
=  75.6. 

Lithium,  6.94  ;   Sodium,  23.00  ;  Potassium,  39.10. 

6.94  + 39.10  =  23 

2 

In  1875  Lenssen  attempted  to  arrange  all  the  elements  known 
in  such  groups  of  three. 

z  337 


338  OUTLINES   OF  CHEMISTRY 

In  1864  Newlands  pointed  out  a  relation  which  he  termed 
the  law  of  octaves.  He  arranged  the  elements  in  the  order  of 
the  magnitude  of  their  atomic  weights,  and  thus  found  that  the 
eighth  element  has  properties  similar  to  the  first,  no  matter 
from  which  element  we  begin  to  count.  He  did  not  work  out 
a  complete  classification,  however. 

In  1869  Dimitri  Mendeleeff,  and  practically  simultaneously 
Lothar  Meyer,  arranged  all  of  the  elements  in  a  table  which 
is  known  as  the  periodic  system  of  the  elements.  In  slightly 
modified  form,  this  table  has  remained  the  best  classification 
of  the  elements  to  the  present  day.  Referring  to  the  table 
on  page  339  it  will  be  noted  that  the  elements  are  arranged 
horizontally  in  the  order  of  magnitude  of  their  atomic  weights, 
in  nine  groups,  which  are  numbered  from  zero  to  VIII.  The 
symbols  in  the  horizontal  series  are  so  written  that  similar 
elements  appear  in  the  same  vertical  column. 

At  the  head  of  each  column  is  indicated  the  valence  of  the 
elements  towards  oxygen  and  also  towards  hydrogen.  So  from 
left  to  right  the  maximum  valence  of  the  elements  towards 
oxygen  increases  from  zero  to  eight,  while  the  valence  towards 
hydrogen  is  greatest  in  group  IV,  i.e.  in  the  middle  of  the 
table. 

Beginning  with  lithium  and  passing  horizontally  to  fluorine, 
the  elements  show  a  gradation  from  powerfully  basic  to  strongly 
acidic  properties.  The  second  horizontal  series,  beginning  with 
sodium  and  ending  with  chlorine,  also  shows  the  same  phenom- 
enon. The  first  and  second  series  are  consequently  two  com- 
plete short  periods.  In  the  third  horizontal  series,  beginning 
with  potassium  and  ending  with  manganese,  iron,  nickel,  and 
cobalt,  we  do  not  have  a  complete  change  from  strongly  basic  to 
strongly  acidic  properties;  but  by  continuing  in  the  fourth 
horizontal  series  from  copper  to  the  right  we  do  pass  from  the 
more  basic  elements  to  the  acidic  bromine.  Consequently  the 
third  and  fourth  horizontal  series  together  are  said  to  form  one 
long  period.  Again,  taking  the  fifth  series,  beginning  with  the 
basic  element  rubidium  and  ending  with  the  metals  ruthenium, 
rhodium,  and  palladium,  and  then  continuing  in  the  sixth  series 
from  silver  to  iodine,  we  have  a  complete  change  from  strongly 
basic  to  markedly  acidic  properties.  For  this  reason,  the  fifth 
and  sixth  horizontal  series  together  are  a  second  long  period.' 


CLASSIFICATION   OE  THE   ELEMENTS 


I    hH 


£X 


OS 


MM 


MM 


eaiaarg 


saoiaaj 


68 


SI 


*s 

32 


Is 


JS 


339 


340  OUTLINES   OF   CHEMISTRY 

The  seventh,  eighth,  ninth,  and  tenth  horizontal  series  taken 
together  are  sometimes  considered  as  the  fifth  period,  and  the 
eleventh  series  as  a  sixth  period.  These  latter  periods,  to  be 
sure,  are  incomplete. 

When  helium,  neon,  argon,  krypton,  and  xenon  were  discov- 
ered, the  question  as  to  their  place  in  the  periodic  system  arose. 
As  these  gases  do  not  combine  with  anything,  their  valence  is 
zero,  and  so  these  elements  have  been  placed  in  a  zero  group  at 
the  head  of  the  system. 

It  will  be  noted  that  the  periodic  system  contains  many  blank 
spaces.  These  represent  elements  yet  to  be  discovered.  When 
Mendeleeff  first  published  his  table,  the  spaces  now  occupied  by 
the  elements  scandium,  gallium,  and  germanium  were  vacant. 
He  boldly  predicted  that  these  metals  would  be  found;  and 
from  the  known  characteristics  of  the  neighboring  elements  in 
the  table,  he  foretold  the  approximate  atomic  weights  and  also 
described  in  some  detail  what  the  physical  and  chemical  prop- 
erties of  these  metals  would  be.  He  called  the  elements 
ekaboron,  ekaaluminium,  and  ekasilicon.  When  some  years 
later  scandium,  gallium,  and  germanium  were  discovered,  their 
atomic  weights  and  other  properties  proved  to  be  those  of  the 
metals  foretold  by  Mendeleeff.  This  brilliant  achievement  of 
the  great  Russian  chemist  attracted  special  attention  to  the 
value  of  the  periodic  system,  which  has  since  served  to  stimu- 
late inquiry  in  various  lines. 

Some  of  the  elements  did  not  seem  to  fit  properly  into  the 
table  and  so  the  question  arose  whether  their  atomic  weights 
had  really  been  correctly  determined.  This  led  to  more  accu- 
rate atomic  weight  determinations  of  a  number  of  elements. 

It  will  be  noted  that  according  to  the  size  of  their  atomic 
weights  tellurium  and  iodine  ought  to  change  places  in  the 
table ;  and  this  ought  also  to  occur  in  the  case  of  argon  and 
potassium.  Considering  the  properties  of  these  elements,  how- 
ever, such  changes  are  not  to  be  thought  of  for  a  moment,  for  it 
would  take  iodine  from  the  column  of  the  halogens,  in  which  it 
certainly  belongs,  and  place  it  in  the  sixth  group  with  oxygen 
and  sulphur.  Similarly,  potassium  must  remain  in  group  I 
with  the  other  alkali  metals.  Redeterminations  of  the  atomic 
weight  of  tellurium  have  shown  that  this  element,  indeed,  has  a 
slightly  higher  atomic  weight  than  iodine  ;  and  so  the  anomalies 


CLASSIFICATION  OF  THE  ELEMENTS 


341 


m 


<M 


§ 


00 


N 


O 

CD 


342  OUTLINES   OF   CHEMISTRY 

mentioned  remain  unexplained.  The  position  of  hydrogen 
in  the  system  is  also  uncertain.  This  element  does  not  seem 
to  fit  into  the  table.  Furthermore,  group  VIII  is  peculiar  as 
compared  with  the  other  groups.  It  contains  three  groups  of 
three  elements  each,  though  to  be  sure  the  elements  in  each  of 
these  groups  have  approximately  the  same  atomic  weight.  A 
number  of  the  rare-earth  elements  (which  see)  have  atomic 
weights  that  are  not  widely  different  from  one  another,  and  for 
these  there  do  not  appear  to  be  suitable  places  in  the  system. 

Stated  in  words,  the  so-called  periodic  law  is  that  the  physical 
and  chemical  properties  of  the  elements  are  periodic  functions  of 
their  atomic  weights.  An  illustration  of  this  is  given  in  Fig.  122, 
in  which  the  atomic  weights  are  represented  as  abscissas  and 
the  atomic  volumes  (i.e.  the  atomic  weights  divided  by  the 
specific  gravities)  are  represented  as  ordinates.  The  trend  of 
the  curve  shows  the  periodicity,  similar  elements  appearing  in 
similar  positions  on  the  curve,  which  was  first  published  by 
Lothar  Meyer. 

The  periodic  system  does  not  represent  a  sharp  quantitative 
relationship  between  the  atomic  weights  and  the  properties  of 
the  elements.  Some  of  its  anomalies  have  already  been  men- 
tioned. In  spite  of  these  imperfections,  the  periodic  system 
offers  a  useful  means  of  classifying  the  elements,  which  will  in 
our  further  considerations  be  grouped  accordingly.  It  will  be 
noted  that  in  considering  the  non-metals  the  natural  families 
that  have  been  studied  really  represent  the  essential  elements 
of  certain  groups  of  the  periodic  system.  The  reader  will 
comprehend  the  significance  of  the  periodic  system  much  better 
after  having  studied  the  physical  and  chemical  peculiarities  of 
the  metals  which  still  remain  to  be  considered. 


CHAPTER  XXI 

THE   ALKALI  METALS 

THE  alkali  metals  are  potassium,  sodium,  lithium,  rubidium, 
and  caesium.  Of  these  potassium  and  sodium  are  by  far  the 
most  abundant  and  important.  Lithium  is  also  found  in  fair 
quantities,  but  rubidium  and  caesium  occur  only  in  very  small 
amounts,  and  they  will  consequently  receive  less  consideration 
here.  None  of  the  metals  of  this  group  are  found  in  the  free 
state  in  nature.  They  always  occur  in  the  form  of  salts,  which 
is  due  to  the  fact  that  their  affinity  for  oxygen,  the  halogens, 
sulphur,  and  other  non-metals  is  very  great.  The  compounds 
which  the  alkali  metals  form  with  the  non-metals  are  in  general 
simple,  very  stable,  and  well  characterized.  These  metals  are 
univalent  in  all  of  their  salts,  hydroxides,  and  oxides.  The 
hydroxides  of  the  alkali  metals  are  the  most  powerful  bases  known. 
The  solutions  of  the  hydroxides  are  very  alkaline  and  caustic, 
so  that  they  are  commonly  called  the  caustic  alkalies.  These 
alkalies  are  non-volatile,  and  are  consequently  termed  the  fixed 
alkalies,  in  contradistinction  to  ammonium  and  other  very  basic 
groups  which  may  under  proper  conditions  be  volatilized.  So 
far  as  its  general  chemical  behavior  is  concerned,  ammonium 
(NH4)  is  closely  allied  to  the  alkali  metals,  and  it  will  conse- 
quently be  advantageous  to  refer  to  the  chemistry  of  the 
ammonium  compounds  in  this  chapter. 

Occurrence,  Preparation,  and  Properties  of  Potassium. — Potas- 
sium is  very  widely  distributed  as  a  constituent  of  silicates  like 
potassium  feldspar  and  certain  forms  of  mica.  Inasmuch  as 
soils  are  produced  by  the  disintegration  of  rocks  by  the  process 
of  weathering,  all  soils  contain  potassium  in  the  feldspathic 
constituents  they  have  derived  from  rocks.  Plants  take  up 
potassium  salts  from  the  soil,  and  in  the  ashes  of  plants  the 
potassium  is  found  as  potassium  carbonate,  commonly  called 
potash.  By  treating  ashes  with  water  and  filtering,  the  potas- 

343 


344  OUTLINES   OF  CHEMISTRY 

slum  carbonate  is  obtained  by  evaporating  the  nitrate.  Since 
human  beings  and  animals  get  their  food  supply  directly  or 
indirectly  from  plants,  it  is  not  strange  that  potassium  is  found 
in  all  animal  tissues  and  secretions,  like  muscles,  bones,  blood, 
urine,  albumen,  eggs,  milk,  etc.  Oceanic  water  contains  about 
0.04  per  cent  potassium,  while  the  earth's  crust  contains  about 
2.45  per  cent  of  the  element.  Though  potassium  is  thus  widely 
distributed,  it  occurs  in  large  quantities  in  but  few  places. 
The  chief  deposits  of  potassium  salts  are  found  in  Germany, 
notably  at  Stassfurt,  where  they  occur  as  layers  twenty  to  thirty 
meters  thick,  covering  strata  of  native  common  salt.  Potassium 
occurs  here  mainly  as  carnallite  KC1  •  MgCl2  •  6  H2O  and  kainite 
MgSO4  -  KC1  •  3  H2O,  but  also  as  sylvite  KC1.  Associated  with 
sodium  nitrate,  potassium  nitrate  is  found  to  some  extent  in 
Peru  and  Chili. 

Metallic  potassium  was  first  prepared  by  Sir  Humphry  Davy, 
who  in  1807  electrolyzed  molten  caustic  potash.  At  present  it 
is  prepared  commercially  by  electrolysis  of  either  potassium  chlo- 
ride or  potassium  hydroxide,  though  formerly  it  was  largely 
made  by  heating  potassium  carbonate  with  carbon:  — 

K2CO3  +  2C  =  2K  +  3CO. 

In  this  process  the  potassium  passes  off  as  vapor  which  is  con- 
densed and  kept  under  petroleum  oils. 

Potassium  is  a  silvery  white  metal,  which  has  a  bright  metal- 
lic luster,  and  is  soft  as  wax  at  ordinary  temperatures.  Below 
0°  it  becomes  hard  and  brittle.  Its  specific  gravity  is  0.865  at 
15°.  It  melts  at  62.5°  and  boils  at  667°.  Its  vapors  are  green. 
The  atomic  weight  of  potassium  is  39.10;  and  its  molecular 
weight  is  the  same,  the  vapors  being  about  twenty  times  as 
heavy  as  hydrogen.  Potassium  reacts  vigorously  with  water, 
evolving  hydrogen  and  forming  potassium  hydroxide.  The  heat 
generated  during  the  action  is  generally  so  great  as  to  set  the 
hydrogen  and  some  of  the  potassium  on  fire,  thus  giving  rise  to 
explosions.  A  freshly  cut  surface  of  potassium  at  once  becomes 
blurred  because  of  reaction  with  the  moisture  of  the  air.  The 
metal  is  consequently  kept  under  petroleum  oils. 

Potassium  hydride  KH  is  formed  by  passing  hydrogen  over 
potassium  at  360°.  It  consists  of  white  needlelike  crystals, 
that  catch  fire  on  exposure  to  the  air.  Water  decomposes 


THE   ALKALI   METALS  345 

the  compound,  forming   potassium   hydroxide  and  hydrogen, 
thus  :  — 

2  KH  +  2  H2O  =  2  KOH  +  2  H2. 

With  carbon  dioxide  it  readily  forms  potassium  formate  :  — 

=  HCOOK. 


Compounds  of  Potassium  with  the  Halogens.  —  Of  these  salts 
potassium  chloride  KC1  is  the  commonest.  In  nature  it  occurs 
as  sylvite  KC1,  and  also  in  carnallite  MgCl2  •  KC1  •  6  H2O,  as 
already  stated.  Potassium  chloride  crystallizes  in  cubes  which 
melt  at  730°  ;  at  higher  temperatures  it  is  converted  into  vapor. 
With  many  other  salts  it  unites  to  form  double  salts,  examples 
of  which  we  have  in  carnallite  and  kainite.  Water  dissolves 
potassium  chloride  readily.  The  salt  is  not  soluble  in  liquid 
hydrochloric  acid  ;  hence  the  addition  of  hydrochloric  acid  to 
a  concentrated  aqueous  solution  of  potassium  chloride  causes 
a  precipitate  of  the  latter  to  form. 

Potassium  bromide  KBr  is  made  by  the  action  of  bromine  on 
potassium  hydroxide  :  — 

6  KOH  +  3  Br2  =  KBrO3  +  5  KBr  +  3  H2O. 

The  potassium  bromate  simultaneously  formed  is  reduced  by 
heating  the  product  with  carbon.  Potassium  bromide  forms 
cubical  crystals  that  melt  at  715°.  It  is  used  in  medicine  and 
in  the  process  of  preparing  silver  bromide  for  photographic 
plates. 

Potassium  iodide  KI  is  prepared  by  the  action  of  iodine  upon 
potassium  hydroxide,  the  process  being  analogous  to  that  de- 
scribed for  making  the  bromide.  The  following  method  is  also 
used  for  making  the  iodide  :  Iodine  is  mixed  with  iron  filings 
under  water,  when  a  solution  of  a  compound  Fe3T8  (that  is, 
FeI2  +  2  FeI3)  is  formed.  This  when  treated  with  potassium 
carbonate  yields  potassium  iodide,  which  remains  dissolved, 
and  an  hydroxide  of  iron  which  is  insoluble  and  can  be  filtered 
off.  Carbon  dioxide  .gas  is  also  given  off  during  the  change, 
which  is  :  — 

Fe3I8  +  4  K2CO3  +  4  H2O  =  8  KI  +  Fe3(OH)8  +  4  CO2. 

Potassium  iodide  crystallizes  in  cubes,  on  evaporation  of  the 
filtrate.  The  salt  melts  at  625°.  It  is  more  copiously  soluble 


346  OUTLINES   OF  CHEMISTRY 

in  water  than  the  bromide.  Its  aqueous  solutions  readily  ac- 
quire a  yellow  color,  due  to  the  separation  of  free  iodine  formed 
by  the  action  of  oxygen  and  carbon  dioxide  of  the  air  upon  the 
salt.  Solutions  of  potassium  iodide  readily  dissolve  additional 
iodine.  These  solutions  are  frequently  employed  in  analytical 
chemistry.  Potassium  iodide  is  often  used  in  medicine,  also 
in  photography. 

Potassium  fluoride  KF  is  formed  by  treating  potassium 
hydroxide  or  carbonate  with  hydrofluoric  acid.  With  the 
latter  it  readily  forms  the  double  compound  KF  •  HF.  Potas- 
sium fluoride  is  a  deliquescent  white  salt,  forming  cubical 
crystals  of  the  composition  KF  +  2  H2O.  The  solutions  attack 
glass. 

All  of  the  halides  of  the  alkali  metals  form  double  salts  with 
salts  of  many  other  metals. 

The  halides  of  the  alkalies  may  all  be  prepared  by  the  action 
of  the  caustic  alkalies  upon  the  hydrohalogen  acids,  or  by  the 
direct  union  of  the  metals  with  the  free  halogens.  The  meth- 
ods employed  in  preparing  the  bromide  and  iodide  of  potassium, 
as  above  described,  are  used  because  of  the  difficulty  of  making 
pure  hydrobromic  and  hydriodic  acids,  to  which  fact  attention 
has  already  been  called. 

Potassium  Hydroxide  KOH.  —  This  compound  is  also  called 
caustic  potash  and  potassium  hydrate.  It  is  prepared  by  treat- 
ing potassium  carbonate  with  slaked  lime  in  vessels  of  iron  or 
silver,  for  caustic  alkalies  attack  glass  or  porcelain.  The 
reaction  is  :  — 

K2CO3  +  Ca(OH)2  =  CaCO3  -f  2  KOH. 

The  calcium  carbonate  is  insoluble,  which  fact  really  forms  the 
basis  of  the  process.  On  evaporating  the  clear  filtrate,  caustic 
potash  is  obtained.  The  latter  is  also  made  in  large  quantities 
by  the  electrolysis  of  solutions  of  potassium  chloride.  In  this 
process  the  electric  current  enters  the  solution  by  a  carbon 
plate  dipping  into  it,  and  leaves  the  solution  by  a  mercury  sur- 
face also  submerged  in  it,  but  not  in  contact  with  the  carbon. 
Thus  as  the  current  passes,  chlorine  is  liberated  on  the  carbon 
and  is  conducted  off  in  pipes  and  used  for  making  bleaching 
powder ;  at  the  same  t^me,  potassium  is  liberated  on  the  mer- 
cury, in  which  it  dissolves.  This  solution  of  potassium  in 


THE   ALKALI   METALS  347 

mercury  is  called  potassium  amalgam.  Water  acts  on  it  slowly, 
forming  potassium  hydroxide  and  hydrogen,  leaving  the  mer- 
cury behind.  The  aqueous  solution  on  evaporation  yields  solid 
caustic  potash.  The  essential  reactions  of  the  process  are,  first, 

by  electrolysis, 

2KC1  =  2K  +  C12; 

and  when  the  amalgam  is  acted  upon  by  water, 
2  K  +  2  H20  =  2  KOH  +  H2. 

Potassium  hydroxide  is  a  hard,  brittle,  white  solid,  which  is 
deliquescent  and  very  soluble  in  water  with  evolution  of  heat. 
The  solution  is  very  caustic,  having  a  corrosive  and  disinte- 
grating action  upon  animal  and  vegetable  tissues.  It  is  the 
most  powerful  of  the  ordinary  bases,  and  consequently  gen- 
erally decomposes  the  salts  of  other  bases.  Caustic  potash 
commonly  comes  into  the  market  cast  in  sticks  which  contain 
about  80  per  cent  of  the  compound  KOH  and  20  per  cent 
water.  Caustic  potash  readily  absorbs  carbon  dioxide,  forming 
potassium  carbonate.  As  a  drying  agent  and  an  absorbent  for 
carbon  dioxide,  potassium  hydroxide  is  much  used  in  chemical 
laboratories,  though  sodium  hydroxide,  which  is  cheaper,  is 
often  employed  in  its  place  when  it  will  do  just  as  well. 
Caustic  potash  is  used  in  making  soft  soaps. 

Potassium  Oxide  K2O  may  be  prepared  (1)  by  melting  potas- 
sium and  potassium  hydroxide  together,  or  (2)  by  heating 
potassium  nitrate  with  potassium  ;  the  reactions  are  :  — 

(1)  2  KOH  +  2  K  =  2.K20  +  H2. 

(2)  2  KNO3  +  10  K  =  6  K2O  +  Na. 

The  oxide  is  a  white,  unstable  powder.  With  water  it  unites, 
yielding  potassium  hydroxide.  Exposed  to  the  air,  it  absorbs 
oxygen,  forming  potassium  peroxide  KO2,  which  is  a  yellow 
powder.  With  water  this  yields  oxygen,  hydrogen  peroxide, 
and  potassium  hydroxide  :  — 

4  K02  +.  6  H20  =  4  KOH  +  4  H2O2  +  Oa. 

Peroxide  of  potassium  is  also  formed  together  with  the  oxide 
when  potassium  is  burned  in  the  air  or  in  oxygen. 

Potassium  Chlorate  KC1O3  may  be  obtained  by  passing 
chlorine  into  a  hot  solution  of  potassium  hydroxide,  as  already 


348  OUTLINES  OF   CHEMISTRY 

described.  By  electrolyzing  a  solution  of  potassium  chloride, 
chlorine  and  potassium  hydroxide  form  at  the  opposite  elec- 
trodes ;  and  by  stirring  the  hot  solution,  the  chlorate  thus  forms 
and  crystallizes  out.  Often  solutions  of  the  more  soluble 
sodium  chlorate  are  thus  prepared,  by  means  of  which  potassium 
chlorate  is  precipitated  from  potassium  chloride  solutions. 
Nearly  all  the  potassium  chlorate  of  commerce  is  now  made  dec- 
trolytically . 

Potassium  chlorate  crystallizes  in  the  monoclinic  system. 
About  six  parts  of  it  dissolve  in  100  parts  of  water  at  room 
temperature.  The  salt  melts  at  350°  and  yields  oxygen  at  a 
slightly  higher  temperature.  It  is  used  for  making  oxygen, 
also  in  manufacturing  matches,  fireworks,  and  explosives.  The 
ease  with  which  the  salt  gives  up  oxygen  is  shown  by  mixing 
two  or  three  grains  of  it  with  a  grain  of  sulphur  or  of  red 
phosphorus  in  a  mortar.  As  the  substances  are  pressed  to- 
gether by  means  of  the  pestle,  there  is  an  explosion.  By  heating 
the  chlorate,  potassium  perchlorate  KC1O4  is  produced  as  a  first 
decomposition  product,  thus  :  — 

8  KC103  =  5  KC104  +  3  KC1  +  2  O2. 

It  forms  rhombic  crystals  and  is  less  soluble  than  either  the 
chlorate  or  chloride,  and  consequently  it  may  readily  be  sepa- 
rated from  these  by  fractional  crystallization.  At  400°  the 
perchlorate  decomposes  into  chloride  and  oxygen : 

KC1O4  =  KC1  +  2  O2. 

Potassium  bromate  KBrO3  and  potassium  iodate  KIO3  are 

analogous  to  the  chlorate.  The  methods  of  their  preparation 
have  already  been  mentioned  under  potassium  bromide  and 
iodide. 

Potassium  Nitrate  KNO3,  also  called  saltpeter,  is  widely  dis- 
tributed in  soils  in  small  quantities,  being  formed  wherever 
organic  substances  decay.  It  was  formerly  produced  on  a 
large  scale  by  allowing  refuse  of  nitrogenous  organic  bodies 
to  decay  in  presence  of  potassium  salts.  At  present  potassium 
nitrate  is  made  by  treating  hot,  saturated  solutions  of  Chili 
saltpeter  NaNO3  with  potassium  chloride,  thus :  — 

NaNO,  +  KC1  =  KNO,  +  NaCl. 


THE   ALKALI  METALS  349 

The  sodium  chloride  formed,  being  far  less  soluble  than  potas- 
sium nitrate,  is  precipitated,  and  from  the  clear  supernatant 
solution  potassium  nitrate  is  readily  obtained  in  form  of  crys- 
tals on  cooling.  The  product  is  further  purified  by  recrystal- 
lizing.  At  0°,  100  parts  of  water  dissolve  13  parts  of  KNO3, 
while  at  100°,  247  parts  of  the  salt  are  dissolved. 

Potassium  nitrate  crystallizes  in  rhombic  prisms,  which 
change  into  rhombohedra  at  about  the  melting  point  of  the 
salt,  339°.  When  heated  above  its  melting  point,  potassium 
nitrate  gives  off  oxygen,  forming  potassium  nitrite  KNO2.  The 
latter  salt  is  more  readily  formed  by  heating  the  nitrate  with 
lead  or  iron,  which  take  up  the  oxygen,  forming  oxides. 

Potassium  nitrate  is  used  as  a  fertilizer,  as  a  preservative  for 
meat,  as  an  oxidizing  agent  in  the  laboratory,  and  as  an  ingredi- 
ent of  fireworks  and  gunpowder.  Most  of  it  is  used  in  making 
gunpowder,  which  consists  of  a  mixture  of  75  parts  saltpeter, 
13  parts  charcoal,  and  12  parts  sulphur.  When  it  is  ignited, 
the  following  reaction  takes  place :  — 

2  KNO3  +  S  +  3C  =  K2S  +  3  CO2+  N2. 

Thus  a  large  volume  of  gas  is  suddenly  liberated,  and  this 
causes  the  explosion.  The  pressure  of  the  gases  produced  at 
2200°,  the  temperature  of  the  discharge,  is  about  96,000  pounds 
per  square  inch.  When  discharged  under  pressure,  as  in  a 
gun,  the  chemical  reaction  is  somewhat  more  complicated  than 
the  one  above  given,  potassium  carbonate,  sulphate,  and  thiosul- 
phate  being  also  produced  in  notable  amounts.  Besides  this, 
products  of  partial  combustion  remain  suspended  in  the  air  in 
a  finely  divided  state,  as  smoke,  after  the  discharge.  Black 
powder,  as  it  is  called,  is  being  displaced  more  and  more  by 
smokeless  powder,  which  see. 

Potassium  Cyanide  KCN  has  already  been  mentioned  in  con- 
nection with  cyanogen.  Besides  being  made  as  there  described, 
it  is  prepared  on  a  large  scale  by  heating  potassium  f  errocyanide 
K4Fe(CN)6  with  potassium  carbonate  :  — 

K4Fe(CN)6  +  K2CO3  =  5  KCN  +  KCNO  +  Fe  +  CO2. 

The  potassium  cyanide  thus  prepared  always  contains  some 
potassium  cyanate  KCNO,  which,  in  case  a  pure  product  is 
required,  is  reduced  by  means  of  charcoal  or  zinc.  On  evapo- 


350  OUTLINES   OF   CHEMISTRY 

< 

ration  of  the  clear  solution,  potassium  cyanide  is  obtained 
in  white  deliquescent  lumps  that  readily  dissolve  in  water. 
Moisture  decomposes  the  salt  somewhat,  yielding  caustic  potash 
and  hydrocyanic  acid.  In  presence  of  carbon  dioxide,  potas- 
sium carbonate  is  formed  together  with  hydrocyanic  acid. 
Hence  the  odor  of  the  latter  is  ever  present  with  potassium 
cyanide.  This  salt  is  extremely  poisonous.  It  is  a  powerful 
reducing  agent,  readily  passing  over  into  the  cyanate,  a  white 
salt  which  also  dissolves  in  water.  Potassium  cyanide  also 
unites  with  sulphur  to  form  potassium  sulphocyanate  KCNS, 
as  already  mentioned.  This  salt  is  crystalline,  deliquescent, 
and  consequently  readily  soluble  in  water. 

Potassium  cyanide  is  used  in  very  large  quantities  for  extract- 
ing gold  from  its  ores.  It  is  also  employed  in  photography  and 
in  gold  and  silver  electroplating,  which  see. 

Potassium  Carbonate  K2CO3  was  formerly  prepared  as  potash 
by  leaching  out  wood  ashes.  The  molasses  residues  of  the  beet 
sugar  industry  and  the  fat  of  sheep's  wool  also  contain  potas- 
sium salts,  from  which  potash  is  obtained.  The  bulk  of  the  po- 
tassium carbonate  on  the  market  is  made  from  potassium  sulphate 
and  chloride  obtained  from  the  Stassfurt  deposits.  The  method 
employed  is  analogous  to  the  Le  Blanc  process  for  making 
sodium  carbonate  (which  see).  In  addition,  however,  potas- 
sium carbonate  is  now  made  at  Neustassfurt  by  passing  carbon 
dioxide  into  magnesium  carbonate  MgCO3  •  3  H2O  suspended 
in  a  solution  of  potassium  chloride,  thus  :  — 

2  KC1  +  3(MgCO3  •  3  H2O)  +  CO2 

=  2(MgCO3  .  KHCOg  -  4  H2O)  +  MgCl2. 

When  the  double  carbonate  thus  formed  is  properly  heated, 
magnesium  carbonate  and  potassium  carbonate  are  obtained  ; 
the  former,  being  insoluble,  is  filtered  off,  and  the  filtrate,  upon 
evaporation,  yields  potassium  carbonate.  The  latter  melts  at 
about  840°.  At  0°,  100  parts  of  water  dissolve  83  parts  of  the 
salt,  while  at  20°,  112  parts  of  the  salt  are  thus  dissolved. 
From  concentrated  solutions,  monoclinic  crystals,  2  K2CO3  4- 

3  H2O,  may  be  obtained.     The  solutions  have  a  strongly  alka- 
line reaction,  for  the  salt  is  one  of  a  weak  acid  with  a  powerful 
base,  and  is  consequently  appreciably  hydrolyzed.     In  the  mar- 
ket, potassium  carbonate  is  often  called  pearlash.     It  is  used 


THE   ALKALI  METALS  351 

in  making  hard  glass  (potash  glass),  soft  soap,  and  many  salts 
of  potassium. 

By  passing  carbon  dioxide  into  solutions  of  potassium  car- 
bonate, potassium  bicarbonate  KHCO3  is  formed.  It  is  less 
soluble  than  the  carbonate,  but  more  so  than  sodium  bicarbon- 
ate, and  hence  cannot,  like  the  latter,  be  obtained  by  the  Solvay 
process  (which  see).  Heating  potassium  bicarbonate,  even  in 
aqueous  solutions,  decomposes  it,  yielding  the  carbonate,  carbon 
dioxide,  and  water. 

Potassium  Silicate  K2SiO3  is  made  by  fusing  silica  with  the 

carbonate : — 

K2CO3  +  SiO2  =  K2SiO8  +  CO2. 

Thus  a  glassy,  deliquescent  mass  is  obtained,  which  dissolves 
in  water,  yielding  a  thick  sirupy  solution  popularly  called 
potassium  water  glass.  The  solutions  commonly  contain  other 
potassium  silicates  besides  K2SiO3.  Its  uses  are  the  same  as 
those  of  the  cheaper  sodium  water  glass  (which  see). 

Potassium  Fluosilicate  K2SiF6  is  formed  as  an  amorphous, 
translucent  precipitate,  when  solutions  of  a  potassium  salt  are 
treated  with  hydrofluosilicic  acid.  It  is  soluble  in  about  800 
parts  of  water  at  20°. 

Potassium  Phosphates.  — The  existence  of  three  phosphates, 
the  primary,  KH2PO4,  the  secondary,  K2HPO4,  and  the  tertiary, 
K3PO4,  has  already  been  mentioned.  They  are  white  salts, 
readily  soluble  in  water.  Their  general  characteristics  have 
been  sufficiently  described. 

Potassium  Sulphate  K2SO4  is  found  at  Stassfurt  in  schonite 
K2SO4  •  MgSO4  •  3  H2O,  from  which  it  is  obtained  by  treatment 
with  potassium  chloride  in  solutions,  thus  :  — 

K2S04  -  MgSO4  +  2  KC1  =  2  K2SO4  +  MgCl2. 

The  magnesium  chloride  is  very  soluble  and  hence  remains  in 
solution,  while  the  less  soluble  potassium  sulphate  is  precipi- 
tated. Potassium  sulphate  is  also  obtained  by  the  action  of 
sulphuric  acid  on  potassium  chloride.  It  crystallizes  in  rhombic 
forms  without  water  of  crystallization.  Its  melting  point  is 
1080°.  At  room  temperature  10  parts  of  the  salt  dissolve  in 
100  parts  of  water.  It  is  used  as  a  fertilizer,  also  in  making 
potassium  alum,  hard  glass,  and  potassium  carbonate. 


352  OUTLINES  OF   CHEMISTRY 

On  treating  potassium  sulphate  with  sulphuric  acid,  acid 
potassium  sulphate  KHSO4  is  formed.  This  is  very  soluble 
in  water  and  melts  at  about  200°,  forming  water  and  potassium 
pyrosulphate  K2S2O7,  which,  on  further  heating,  decomposes 
into  the  normal  sulphate  and  sulphur  trioxide.  The  latter  at 
higher  temperatures  readily  unites  with  many  metallic  oxides ; 
hence  the  practice  of  fusing  refractory  metallic  oxides  and  many 
minerals  with  potassium  bisulphate  KHSO4  to  convert  the  bases 
into  soluble  sulphates. 

Potassium  Sulphite  K2SO3  is  formed  by  passing  sulphur 
dioxide  into  a  solution  of  potassium  carbonate  till  carbon  diox- 
ide is  no  longer  formed.  The  salt  crystallizes  in  monoclinic 
prisms  with  two  molecules  of  water. 

On  saturating  a  solution  of  potassium  carbonate  or  sulphite 
with  sulphur  dioxide,  acid  potassium  sulphite  or  potassium  bisul- 
phite KHSO3,  crystallizing  in  needles,  is  obtained. 

Sulphides  of  Potassium. — Potassium  sulphide  K2S  is  made  by 
fusing  potassium  sulphate  with  charcoal:  — 

K2S04  +  4  C  =  K2S  +  4  CO. 

It  is  a  flesh-colored,  crystalline  mass  that  readily  dissolves  in 
water.  The  oxygen  of  the  air  acts  on  the  solutions,  gradually 
forming  potassium  thiosulphate  K2S2O3,  thus  :  — 

2  K2S  +  H20  +  2  02  =  2  KOH  +  K2S2O3. 

By  saturating  a  caustic  potash  solution  with  hydrogen  sul- 
phide, potassium  sulphydrate  KSH  is  formed  :  — 

KOH  +  H2S  =  KSH  +  H2O. 

Its  solutions  are  alkaline,  and  upon  evaporation  with  caustic 
potash  they  yield  potassium  sulphide:  — 

KSH  +  KOH  =  K2S  +  H2O. 

Solutions  of  K2S  or  KSH  will  readily  dissolve  sulphur,  forming 
a  series  of  compounds  known  as  polysulphides.  Thus  K2S3, 
K2S4,  and  K2S5  have  been  obtained.  On  treatment  with  acids 
the  polysulphides  yield  hydrogen  sulphide  and  sulphur.  By 
fusing  potash  with  sulphur  out  of  contact  with  air,  a  mass  of 
liver-brown  color,  known  as  liver  of  sulphur,  is  obtained.  It 
consists  of  a  mixture  of  polysulphides  of  potassium  together 
with  potassium  sulphate  and  thiosulphate. 


THE   ALKALI  METALS  353 

Tests  for  Potassium.  —  To  detect  the  presence  of  potassium 
in  small  quantities,  the  spectroscope  (which  see)  is  employed. 
In  addition,  the  fact  that  acid  potassium  tartrate  KHC4H4O6, 
potassium  silicofluoride  K2SiF6,  and  potassium  platinic  chloride 
K2PtCl6  are  difficultly  soluble  in  water  serves  to  determine 
whether  potassium  salts  are  present  in  a  given  solution  or  not. 
The  reactions  involved,  written  for  potassium  chloride,  are  as 
follows  :  — 

KC1  +  H2C4H406  =  KH  •  C4H4O6  +  HC1. 

2  KC1  +  H2SiF6  =  K2SiF6  +  2  HC1. 
2  KC1  +  PtCl4  =  K2PtCl6. 

Potassium  platinic  chloride  is  soluble  in  about  100  parts  of 
water;  but  its  solubility  is  much  less  in  alcoholic  solutions,  of 
which  fact  the  analytical  chemist  avails  himself. 

Rubidium  and  Caesium.  —  These  metals  have  the  atomic 
weights  85.45  and  132.81  respectively.  They  were  discovered 
in  1860  by  Robert  Bunsen  by  means  of  the  spectroscope,  in  the 
residues  obtained  by  evaporating  Durkheim  mineral  water. 
The  spectrum  of  rubidium  contains  certain  characteristic  red 
lines,  while  that  of  caesium  exhibits  striking  blue  lines.  Hence 
Bunsen  named  the  elements  rubidium  (red)  and  caesium 
(blue) ;  their  symbols  are  Rb  and  Cs.  Salts  of  both  of  these 
metals  occur  widely  distributed  in  nature,  though  in  extremely 
minute  quantities;  they  are  commonly  associated  with  salts  of 
potassium.  Carnallite  contains  about  0.025  per  cent  of  rubid- 
ium, and  it  has  been  estimated  that  from  the  Stassfurt  salts  that 
are  used  as  fertilizers  more  than  200  tons  of  rubidium  are  dis- 
tributed annually  over  the  soils,  from  which  plants  absorb  it. 
Thus  in  the  ash  of  sugar  beets,  tobacco,  coffee,  and  tea,  rubidium 
is  frequently  met.  The  methods  that  are  used  for  preparing 
potassium  also  serve  for  making  rubidium.  The  chief  source  of 
rubidium  salts  is  the  Stassfurt  deposit.  Caesium  is  much  rarer 
than  rubidium.  The  mineral  pollux  found  on  the  island  of 
Elba  is  a  silicate  of  aluminum  and  caesium.  It  contains  about 
34  per  cent  of  caesium  oxide. 

Metallic  caesium  was  first  prepared  in  1881  by  Setterberg, 
who  electrolyzed  the  cyanide  CsCN.  Caesium  may  also  be  ob- 
tained by  heating  its  oxide  or  carbonate  with  magnesium. 

The  salts  of  both  rubidium  and  caesium  are  in  general  analogous 

2A. 


354  OUTLINES   OF   CHEMISTRY 

to  those  of  potassium,  with  the  exception  that  the  former  elements 
are  also  able  to  form  halides  containing  three  or  five  halogen  atoms. 
In  such  compounds  rubidium  and  caesium  consequently  exhibit 
valences  of  three  and  five,  respectively.  The  hydroxide  of 
rubidium  is  a  stronger  base-  than  that  of  potassium,  while  the 
hydroxide  of  caesium  is  the  most  powerful  base  known. 

Occurrence,  Preparation,  and  Properties  of  Sodium.  —  The 
chief  compound  of  sodium  is  sodium  chloride,  which  is  widely 
distributed  in  nature  in  large  quantities.  Thus,  it  occurs  in 
oceanic  waters,  while  many  salt  seas  and  lakes  are  practically 
saturated  solutions  of  common  salt.  Mineral  springs  are  often 
rich  in  sodium  salts,  which  also  occur  in  huge  deposits  as  chlo- 
ride, nitrate,  and  borate,  in  various  parts  of  the  globe.  Cryolite, 
an  aluminum  sodium  fluoride,  is  found  in  Greenland,  and  albite 
or  soda  feldspar,  a  silicate  of  sodium  and  aluminum,  is  widely 
distributed  in  nature.  Just  as  land  plants  contain  potassium, 
so  sea  plants  contain  sodium,  which  is  found  in  their  ashes  as 
carbonate.  From  the  soil,  sodium  gets  into  plants  and  then  into 
animal  organisms,  where  it  occurs  in  the  blood  and  the  various 
tissues  and  secretions. 

Metallic  sodium  was  first  prepared  by  Sir  Humphry  Davy 
in  1807  by  electrolysis  of  molten  sodium  hydroxide.  In  this 
way  it  is  now  prepared  on  a  large  scale.  Thus,  sodium 
is  deposited  at  one  pole,  and  the  hydroxyl  liberated  at  the 
opposite  pole  at  once  decomposes,  yielding  water  and  oxygen. 
The  yield  is  only  about  40  per  cent,  for  some  of  the  metal  con- 
tinually reacts  with  the  water,  liberating  hydrogen  and  forming 
sodium  hydroxide.  The  methods  described  for  making  potassium 
may  also  be  used  for  preparing  sodium. 

Sodium  has  properties  similar  to  those  of  potassium.  At  room 
temperature  the  metal  is  soft  like  wax,  while  at  —  20°  it  is  hard. 
It  has  a  bright,  silvery  luster.  Its  melting  point  is  95.5°,  and 
its  boiling  point  742°.  The  vapor  of  sodium  is  colorless,  though 
in  thick  layers  it  appears  violet.  The  vapor  is  12  times  as  heavy 
as  hydrogen,  whence  the  molecular  weight  is  24.  The  atomic 
weight  is  23,  and  the  valence  is  one.  The  molecules  of  so- 
dium, like  those  of  potassium,  consist  of  but  one  atom.  Sodium 
decomposes  water,  like  potassium,  but  the  action  is  not  as  violent 
as  in  the  case  of  the  latter.  The  specific  gravity  of  sodium  is 
0.974  at  15°.  The  metal  is  used  in  making  sodium  cyanide  and 


THE   ALKALI   METALS  855 

sodium  peroxide,  also  in  the  manufacture  of  complex  organic 
compounds.  In  the  laboratory,  it  is  frequently  employed  as  a 
reducing  agent.  Like  potassium,  it  is  kept  under  petroleum 
oil ;  but  sodium  is  also  shipped  in  tightly  soldered  tinned  iron 
boxes. 

Sodium  dissolves  in  mercury,  forming  sodium  amalgam,  which 
when  treated  with  water  yields  mercury,  sodium  hydroxide,  and 
hydrogen.  The  latter  is  liberated  much  less  rapidly  than  when 
sodium  alone  acts  on  water,  hence  the  amalgam  is  often  em- 
ployed in  effecting  reductions  that  are  to  proceed  slowly.  With 
potassium,  sodium  forms  alloys,  1  part  sodium  to  2  to  10  of 
potassium,  that  are  liquid  at  room  temperature  and  have  the 
appearance  of  mercury.  If  but  little  potassium  is  alloyed 
with  much  sodium,  that  is,  if  the  proportions  mentioned  are 
reversed,  the  alloys  formed  are  brittle  solids.  With  hydrogen, 
sodium  forms  sodium  hydride  NaH,  which  is  similar  to  the 
hydride  of  potassium. 

Sodium  Chloride  NaCl,  common  salt,  is  the  chief  source  of  sodium 
and  its  compounds.  Large  deposits  are  found  at  Stassfurt  and 
Reichenhall,  in  Germany,  at  Wieliczka  in  Galicia,  at  Cheshire 
in  England,  at  Syracuse  in  New  York,  in  Michigan,  Kansas, 
Texas,  Utah,  California,  as  well  as  in  Asia  and  Africa.  In  some 
localities  the  salt  is  mined  in  solid  form,  and  again  it  is  fre- 
quently obtained  from  brines  by  evaporation.  When  the  brine 
is  dilute,  as  in  case  of  sea  water,  it  is  concentrated  by  allowing 
it  to  trickle  over  a  large  surface  of  twigs,  thus  giving  better  op- 
portunity for  evaporation  to  take  place.  The  concentrated  solu- 
tion thus  obtained  is  then  generally  evaporated  by  means  of  arti- 
ficial heat ;  though  in  hot  climates,  solar  heat  is  often  relied  upon 
entirely.  Brine  is  also  evaporated  in  shallow  basins  either  by 
means  of  artificial  heat  or  the  heat  of  the  sun.  Again,  so-called 
vacuum  pans,  in  which  brines  are  evaporated  under  diminished 
pressure,  are  frequently  employed.  Ordinary  salt  is  not  pure. 
It  contains  small  amounts  of  sodium  sulphate  and  chlorides  of 
calcium  and  magnesium.  The  latter  are  deliquescent,  and  their 
presence  in  common  salt  causes  it  to  attract  moisture  from  the 
air.  Pure  common  salt  is  obtained  by  conducting  hydrochloric 
acid  gas  into  a  saturated  solution  of  common  salt.  The  latter 
is  thus  thrown  out  of  solution,  for  it  is  less  soluble  in  hydro- 
chloric acid  than  in  water. 


356 


OUTLINES  OF  CHEMISTRY 


FIG.  123. 


During  the  year  1910  the  United  States  produced  4,242,793 
tons  of  sodium  chloride,  which  is  approximately  one  fourth  of 
the  world's  annual  output. 

Sodium  chloride  crystallizes  in  cubes,  which  when  obtained 
from  aqueous  solutions  form  groups  of  hollow  pyramids,  that 

is,  are  hopper-shaped  (Fig.  123). 
They  contain  occluded  water  which 
escapes  on  heating,  causing  the  salt 
to  decrepitate.  At  —  10°,  sodium 
chloride  crystallizes  with  two  mole- 
cules of  water,  forming  monoclinic 
prisms.  Sodium  chloride  melts  at 
about  800°,  at  which  temperature  it 
also  volatilizes  appreciably.  At 
room  temperature,  100  parts  of 
water  dissolve  36  parts  of  salt,  and  at  100°,  39  parts.  The 
salt  is  therefore  about  as  soluble  in  hot  as  in  cold  water.  In 
practically  all  solvents  other  than  water,  common  salt  is  insoluble. 
Human  beings  and  many  animals  cannot  live  without  sodium 
chloride,  which  is  found  distributed  throughout  their  systems 
in  small  amounts,  though  its  function  is  not  known.  It  has 
been  estimated  that  the  quantity  of  salt  used  annually  by  a 
i  \iman  being  amounts  to  about  one  tenth  of  his  weight. 

Common  salt  is  used  in  enormous  quantities  in  making 
hydrochloric  acid,  chlorine,  and  nearly  all  of  the  sodium  com- 
pounds. 

Sodium  bromide  NaBr  and  sodium  iodide  Nal  are  more  solu- 
ble in  water  than  the  chloride;  they  are  also  more  volatile.  In 
general,  these  halides  and  sodium  fluoride  are  much  like  the 
corresponding  salts  of  potassium. 

Oxides  and  Hydroxide  of  Sodium.  — When  sodium  is  burned  in 
oxygen,  a  mixture  of  the  two  oxides,  Na2O  and  Na2O2  is  formed. 
Sodium  oxide  Na2O  is  also  obtained  as  a  gray  mass  by  heat- 
ing the  hydroxide  with  sodium.     When  treated  with  water,  it 
dissolves,  forming  a  solution  of  the  hydroxide. 

Sodium  peroxide  Na2O2  is  a  white  powder  formed  by  pro- 
longed heating  of  sodium  or  sodium  oxide  in  oxygen  or  in  the 
air  at  about  300°.  At  high  temperatures  it  decomposes  into 
the  oxide  and  oxygen.  With  water  it  yields  sodium  hydroxide 
and  oxygen,  though  in  the  cold  it  may  in  part  be  dissolved. 


THE   ALKALI  METALS 


357 


With  dilute  acids  it  yields  hydrogen  peroxide ;  hence  it  may 
be  regarded  as  the  latter  substance  whose  hydrogen  atoms  are 
replaced  by  sodium.  Sodium  peroxide  is  prepared  on  a  large 
scale  as  an  oxidizing  and  bleaching  agent.  It  is  shipped  in  air- 
tight tin  cans. 

Sodium  hydroxide  NaOH,  caustic  soda,  is  made  by  the  same 
methods  as  caustic  potash,  namely,  by  treating  the  carbonate 
with  slaked  lime  or  by  electrolysis  of  solutions  of  the  chloride, 
using  a  mercury  cathode.  It  is  also  made  by  the  Acker  process, 
which  consists  of  electrolyzing  fused  sodium  chloride,  using  a 


FIG.  124. 


carbon  anode  and  a  cathode  of  molten  lead.  The  chlorine  is 
conducted  off  and  used  in  making  bleaching  powder,  while  the 
sodium  forms  an  alloy  with  the  lead.  A  jet  of  steam  directed 
on  this  alloy  reacts  with  the  sodium,  forming  hydrogen,  which 
burns  at  once,  and  sodium  hydroxide,  which,  being  molten,  is 
drawn  off  into  proper  containers.  Figure  124  shows  the  essen- 
tial features  of  the  apparatus. 


358 


OUTLINES   OF   CHEMISTRY 


The  properties  of  sodium  hydroxide  are  much  like  those  of 
potassium  hydroxide.  The  former  is  much  cheaper,  however, 
and  is  used  in  place  of  potassium  hydroxide  whenever  possible. 


FIG.  125. 

Large  quantities  of  sodium  hydroxide  are  used  in  making 
soap  and  in  "  softening  water."  It  is  also  used  in  the  paper 
industry  and  in  making  carbolic  acid,  oxalic  acids,  and  many 
other  products. 

Sodium  Carbonate  Na2CO3,  popularly  called  soda  or  sal  soda, 
is  manufactured  from  common  salt  on  a  very  large  scale, 
for  it  is  essential  in  making  glass,  soap,  caustic  soda,  and 
other  sodium  compounds.  Sodium  carbonate  is  found  in 
nature  in  Wyoming,  California,  Mexico,  Egypt,  and  in  the 
ashes  of  marine  plants.  It  is  manufactured  by  the  Le  Blanc 
process,  the  Solvay  process,  and  the  electrolytic  process. 

The  Le  Blanc  soda  process  was  introduced  in  France  by  Le 
Blanc  in  1791.  It  is  based  upon  three  chemical  reactions. 
First,  common  salt  is  treated  with  sulphuric  acid  in  equivalent 
quantity ;  thus  hydrochloric  acid  and  "  salt  cake,"  consisting 
of  sodium  chloride  and  sodium  acid  sulphate,  are  formed  :  — 

2  NaCl  +  H2S04  =  NaCl  +  NaHSO4  +  HCL 

This  salt  cake  is  then  transferred  to  the  hearth  of  a  furnace, 
(Fig,  125)  and  heated  ;  thus  sodium  sulphate  and  more  hydro- 
chloric acid  are  produced :  — 

NaCl  +  NaHS04  =  Na2SO4  +  HCL 

The  hydrochloric  acid  formed  is  absorbed  in  water  and  placed  on 
the  market.  The  second  step  consists  in  mixing  the  sodium  sul- 
phate formed  with  charcoal  and  calcium  carbonate,  and  heating 


THE   ALKALI   METALS 


359 


the  mixture  in  a  rotating  cylindrical  furnace  (Fig.  126)  which 
has  a  central  flue  through  which  the  hot  gases  of  the  furnace 
pass.  Thus  sodium  sulphate  is  reduced  to  sulphide  :  — 

Na2SO4  +  40  =  Na2S  +  4  CO  ; 

and  sodium  sulphide  reacts  with  calcium  carbonate,  yielding 
sodium  carbonate  and  calcium  sulphide  :  — 

Na2S 


CaCO3  =  CaS 


Na2CO3. 


The  third  step  consists  in  treating  the  mixture  of  calcium  sul- 
phide and  sodium  carbonate,  called  black  ash,  with  water.  Thus 
sodium  carbonate  dissolves  and  calcium  sulphide  remains  behind, 
as  an  insoluble  residue.  By  evaporation  of  the  clear  solution 


FIG.  126. 

crystals  of  the  composition  Na2CO3  •  H2O  are  obtained,  which 
are  heated  to  drive  off  the  water,  thus  forming  Na2CO3,  calcined 
soda.  When  this  is  recrystallized  from  water  at  room  temper- 
atures, crystals  of  the  composition  Na2CO3- 10  H2O  are  formed. 
This  is  the  washing  soda  or  crystallized  soda  of  commerce. 
Sodium  carbonate  made  by  the  Le  Blanc  process  generally  con- 
tains small  amounts  of  chloride,  sulphate,  and  sulphite. 

The  sulphur  contained  in  the  calcium  sulphide,  also  called 
tank  waste,  is  recoveredby  a  process  developed  in  1888  by  Chance 
in  England.  It  consists  of  diluting  the  waste  with  water  and 
arranging  it  in  a  series  of  cylinders  through  which  carbon 


360  OUTLINES   OF  CHEMISTRY 

dioxide  from  a  limekiln  is  run.  There  is  first  formed  calcium 
carbonate  and  calcium  sulphydrate,  thus  :  — 

2  CaS  +  CO2  +  H2O  =  CaCO3  +  Ca(SH)2. 

By  further  action  of  carbon  dioxide,  this  sulphydrate  is  decom- 
posed, yielding  hydrogen  sulphide  :  — 

Ca(SH)2  +  CO2  +  H2O  =  CaCO3  4-  2  H2S. 

The  hydrogen  sulphide  set  free  in  one  cylinder  reacts  with  the 
material  in  the  next,  and  so  on  through  the  series  of  cylinders ; 
thus, 

CaS  +  H2S  =  Ca(SH)2. 

Finally,  the  hydrogen  sulphide  is  burned  to  water  and  sulphur, 
thus, 

2  H2S  +  O2  =  2  H2O  +  2  S ; 

or  it  is  burned  completely  to  water  and  sulphur  dioxide,  and 
then  made  into  sulphuric  acid.  About  90  per  cent  of  the  sul- 
phur in  the  tank  wastes  may  be  recovered. 

The  Solvay  process,  also  called  the  ammonia  soda  process,  was 
perfected  by  the  Belgian  chemist  Solvay  in  1863.  It  is  based 
upon  the  fact  that  sodium  bicarbonate  NaHCO3  is  relatively 
sparingly  soluble  in  water.  The  process  consists  of  conducting 
ammonia  and  carbon  dioxide  into  a  cold,  saturated  solution  of 
common  salt.  In  this  way  ammonium  bicarbonate  is  formed, 
which  reacts  with  sodium  chloride,  causing  a  precipitate  of  so- 
dium bicarbonate  to  separate  out :  — 

NaCl  +  NH4HCO3  =  NH4C1  +  NaHCO3. 

The  sodium  bicarbonate  is  placed  on  the  market  as  such,  or  it 
is  heated  and  thus  converted  into  carbonate, 

2  NaHC03  =  Na2C03  +  CO2  +  H2O. 

The  ammonium  chloride  remains  in  solution,  from  which,  by 
heating  with  lime  or  magnesia,  ammonia  is  again  regenerated. 
Thus,  the  waste  products  are  chlorides  of  calcium  or  magnesium. 
On  heating  magnesium  chloride,  magnesium  oxide  and  hydro- 
chloric acid  are  formed  :  — 

MgCl2  -f  H2O  =  MgO  +  2  HCL 


THE   ALKALI  METALS  361 

In  this  way  the  magnesia  can  be  recovered  and  used  over  again. 
In  Germany  and  the  United  States  most  of  the  sodium  carbon- 
ate is  made  by  the  Solvay  process,  while  in  England  the  Le  Blanc 
process  still  predominates.  It  is  evident  that  a  purer  product 
is  obtained  by  the  Solvay  process. 

The  electrolytic  process  consists  of  making  caustic  soda  by 
one  of  the  electrolytic  methods  described  and  then  treating  the 
solution  with  carbon  dioxide. 

At  ordinary  temperatures  sodium  carbonate  crystallizes  in 
monoclinic  prisms  of  the  composition  Na2CO3  •  10  H2O,  which 
effloresce.  At  60°  this  compound  melts  in  its  crystal  water 
and  on  continued  heating  it  yields  a  deposit  of  the  composition 
Na2CO3  •  2  H2O,  which  when  dried  in  the  air  readily  forms 
Na2CO3  •  H2O.  At  100°  the  salt  may  be  completely  dehydrated. 
At  15°,  100  parts  of  water  dissolve  55  parts  of  Na2CO3,  at  38°, 
138  parts,  and  at  100°,  100  parts.  The  solution  deposits 
Na2CO3  •  7  H2O  at  50°.  Solutions  of  sodium  carbonate  have  a 
strong  alkaline  reaction ;  for,  like  potassium  carbonate,  the  salt 
is  hydrolyzed.  Sodium  carbonate  melts  at  red  heat,  forming  a 
clear  liquid. 

Sodium  bicarbonate  NaHCO3,  or  sodium  acid  carbonate,  is 
soluble  in  about  10  parts  of  water  at  20°.  Its  solutions  have 
an  alkaline  reaction,  for  the  salt  is  decomposed  by  hydrolysis 
to  a  slight  extent.  When  warmed,  the  solutions  give  off  car- 
bon dioxide,  the  carbonate  being  formed.  Sodium  bicarbonate 
is  used  in  medicine,  in  baking  powder,  as  saleratus,  and  in  mak- 
ing soda  water  and  other  effervescent  beverages. 

Sodium  Nitrate  NaNO8  occurs  in  large  quantities  in  Chili  and 
Peru  as  Chili  saltpeter,  which  also  contains  other  salts  of  sodium, 
notably  the  iodate,  sulphate,  and  chloride.  The  salt  crystallizes 
in  the  rhombohedral  division  of  the  hexagonal  system  and  melts 
at  318°.  In  general,  its  chemical  behavior  is  like  that  of  potas- 
sium nitrate.  Sodium  nitrate  is  used  as  a  fertilizer,  also  in  the 
manufacture  of  potassium  nitrate  and  nitric  acid.  The  salt  is 
somewhat  hygroscopic,  which  unfits  it  for  use  in  gunpowder. 

By  heating  sodium  nitrate  with  lead  or  iron,  sodium  nitrite 
NaNO2  is  formed.  This  salt  is  much  used  in  the  coal-tar  dye- 
stuff  industry. 

Phosphates  of  Sodium.  —  The  most  important  of  these  is  the 
secondary  or  disodium  phosphate  Na2HPO4  •  12  H2O.  This  is 


362 


OUTLINES   OF  CHEMISTRY 


commonly  called  simply  sodium  phosphate.  It  is  prepared  by 
the  action  of  phosphoric  acid  on  sodium  hydroxide  or  carbonate. 
The  crystals  effloresce.  At  20°,  100  parts  of  water  dissolve  9.3 
parts  of  the  salt.  The  solution  has  a  slightly  alkaline  reaction. 
On  adding  another  equivalent  of  caustic  soda  to  disodium  phos- 
phate solution  and  evaporating  to  dryness,  the  tertiary  sodium 
phosphate  Na3PO4  •  12  H2O  is  obtained.  In  aqueous  solution,  this 
does  not  exist,  being  hydrolyzed  into  disodium  phosphate  and 
sodium  hydroxide.  Primary  sodium  phosphate  NaH2PO4  •  H2O 


60 


50 


s 
b/, 

§40 


30 


20 


~   10 


S  0        10       20       30      40       50       60      70      80       90      100 

FIG.  127. 

O 

Temperature  in  degrees  C. 

has  an  acid  reaction  in  solutions.     On  heating,  the  salt  loses 
water  and  forms  sodium  metaphosphate  NaPO3. 

Sodium  Sulphate  Na2SO4  - 10  H2O,  also  called  Glauber's  salt, 
crystallizes  in  large  monoclinic  crystals  that  effloresce  on  expos- 
ure to  the  air.  It  is  manufactured  as  one  of  the  products  of 
the  Le  Blanc  soda  process,  also  by  Hargreave's  process,  which 
consists  of  passing  air,  sulphur  dioxide,  and  steam  over  sodium 
chloride  at  high  temperatures,  thus  :  — 

2  NaCl  +  H20  +  O  +  S02  =  Na2SO4  +  2  HC1. 

At  Stassfurt  the  salt  is  formed  by  the  action  of  magnesium 
sulphate  upon  sodium  chloride  at  low  temperatures,  when  the 
sodium  sulphate  is  deposited,  while  the  magnesium  chloride 
simultaneously  formed  remains  in  solution:  — 

MgSO4  +  2  NaCl  =  Na2SO4  +  MgCl2. 


THE   ALKALI  METALS  363 

The  salt  was  first  prepared  by  Glauber,  whence  its  name.  The 
solubility  of  Glauber's  salt  Na2SO4  •  10  H2O  increases  with  rise 
of  temperature  to  32.4°,  beyond  which  it  decreases,  for  at  higher 
temperatures  the  salt  loses  water,  becoming  anhydrous  Na2SO4, 
which  is  less  soluble  than  the  decahydrate.  Figure  127  shows 
the  solubility  curve,  which  has  a  sharp  point  of  inflection  at  32.4°. 
This  point  is  really  the  intersection  of  the  solubility  curve  of 
the  decahydrate  and  that  of  the  anl^drous  salt.  Glauber's  salt 
melts  in  its  water  of  crystallization  at  32.4°.  When  completely 
melted,  the  solution  so  obtained  may  be  cooled  to  room  tempera- 
ture and  even  lower.  This  is  a  so-called  supersaturated  solu- 
tion, for  it  contains  more  salt  than  would  be  taken  up  at  the 
lower  temperature  in  presence  of  an  excess  of  the  solid  salt. 
Indeed,  if  a  crystal  of  the  solid  Na2SO4  •  10  H2O  is  introduced 
into  this  supersaturated  solution,  the  whole  of  it  at  once  solidi- 
fies. Many  other  substances  form  similar  supersaturated  solu- 
tions. Glauber's  salt  simply  furnishes  an  excellent  illustration. 
Sodium  sulphate  occurs  in  many  mineral  waters.  As  Glauber's 
salt  it  is  used  as  a  purgative. 

On  treating  sodium  sulphate  with  an  equivalent  quantity  of 
sulphuric  acid,  acid  sodium  sulphate  or  sodium  bisulphate 
NaHSO4  •  H2O  is  obtained.  It  becomes  anhydrous  above  50° 
and  melts  at  about  300°.  Its  behavior  is  similar  to  that  of  potas- 
sium bisulphate. 

Sodium  Sulphite  Na2SO3  •  7  H2O  is  formed  by  passing  sulphur 
dioxide  into  a  concentrated  solution  of  sodium  hydroxide  or 
carbonate  to  neutrality.  On  saturating  a  strong  solution  of  the 
sulphite  with  sulphur  dioxide,  sodium  bisulphite  NaHSO3  is 
formed.  This  is  frequently  used  as  a  source  of  sulphur  dioxide 
in  bleaching  fabrics  of  silk,  wool,  etc. 

Sodium  Thiosulphate  Na2S2O3  •  5  H2O  is  made  by  boiling  sul- 
phur in  a  solution  of  sodium  sulphite.  It  is  also,  though 
wrongly,  called  hyposulphite  of  soda  or  hypo.  Its  solution 
serves  in  photography  as  a  "  fixing  bath,"  for  it  dissolves  the 
excess  of  silver  bromide  on  the  photographic  plate  after  the 
latter  has  been  exposed  to  the  light  and  developed.  Sodium 
thiosulphate  absorbs  chlorine  and  is  used  as  so-called  "  antichlo- 
rine,"  thus:  — 

2  Na2S203  +  C12  =  2  NaCl  +  Na2S4O6. 

The  salt  is  also  used  in  chemical  analysis. 


364  OUTLINES  OF  CHEMISTRY 

The  sulphides  of  sodium  and  sodium  hydrosulphide  are  analo- 
gous to  those  of  potassium  and  need  no  special  description. 

Sodium  Silicate,  or  sodium  water  glass,  is  made  by  fusing  silica 
with  sodium  hydroxide  or  carbonate,  or  with  Glauber's  salt  and 
carbon.  It  is  also  prepared  by  boiling  silica  with  concentrated 
solutions  of  sodium  hydroxide  under  pressure.  Sodium  silicate 
may  be  obtained  in  form  of  monoclinic  crystals  of  the  composi- 
tion Na2SiO3  •  8  H2O.  Water  glass  comes  in  the  market  as  a 
thick  sirupy  solution  containing  various  sodium  silicates,  the 
composition  averaging  about  Na2Si4O9.  It  is  used  in  laundry 
soaps  as  a  "filler,"  in  making  fabrics  and  wood  fireproof,  in  the 
production  of  artificial  stone,  as  a  preservative  for  wood,  and  as 
a  cement  for  glass,  asbestus,  mineral  wool,  etc. 

Sodium  Cyanide  NaCN  is  very  similar  to  potassium  cyanide, 
and  is  used  for  the  same  purposes.  It  is  made  commercially 
by  the  action  of  ammonia  gas  upon  a  mixture  of  carbon  and 
metallic  sodium,  and  is  extremely  soluble  in  water. 

Sodium  Borate  Na2B4O7  •  10  H2O,  or  borax,  has  been  described 
under  boron. 

Lithium  and  its  Compounds.  —  Lithium  is  found  in  the  miner- 
als, lepidolite  or  lithia  mica,  spodumene,  triphylite,  and  a  few 
others  of  rare  occurrence.  In  very  minute  quantities  lithium 
salts  are  also  found  widely  distributed  in  soils,  from  which  they 
pass  into  plants,  a  number,  like  tobacco  and  beets,  being  particu- 
larly prone  to  store  up  lithium.  Many  mineral  waters  contain 
lithium,  though  generally  in  rather  small  amounts.  Lithium 
compounds  are  readily  detected  by  means  of  the  spectroscope; 
for  they  exhibit  two  characteristic  red  lines.  To  the  Bunsen 
flame  lithium  salts  impart  a  characteristic  red  color. 

Metallic  lithium  may  be  obtained  by  electrolyzing  the  fused 
chloride  or  a  concentrated  solution  of  the  latter  in  pyridine. 
It  has  a  very  low  specific  gravity,  0.59,  and  hence  floats  even 
on  petroleum  oils.  In  general,  its  chemical  behavior  is  similar 
to  that  of  sodium,  only  less  vigorous.  The  atomic  weight  of 
lithium  is  6.94,  arid  its  valence  is  1. 

Lithium  chloride  LiCl  is  deliquescent  and  extremely  soluble 
in  water.  On  the  other  hand,  lithium  carbonate  Li2CO3  is 
very  slightly  soluble,  for  100  parts  of  water  dissolve  but  0.77 
part  at  15°.  Lithium  phosphate  Li3PO4  is  also  but  slightly 
soluble,  1  part  in  2500  of  water.  This  slight  solubility  of  the 


THE   ALKALI   METALS 


365 


carbonate  and  phosphate  distinguishes  lithium  sharply  from  the 
other  alkali  metals  and  shows  its  similarity  to  the  alkaline  earth 
metals.  The  chloride  of  lithium  dissolves  in  pyridine,  whereas 
the  chlorides  of  all  the  other  alkali  metals  are  insoluble  in  that 
liquid,  which  fact  is  used  in  separating  lithium  from  the  other 
members  of  the  group. 

Commercially,  lithium  carbonate  is  the  most  important  of  the 
lithium  salts.  With  uric  acid  C5H4N4O3,  lithium  forms  a 
moderately  soluble  salt,  lithium  urate  C5H2N4O3Li2,  upon  which 
fact  is  based  the  administration  of  lithium  carbonate  in  medi- 
cine, in  cases  of  gout,  which  is  caused  by  deposits  of  uric  acid 
in  the  joints  and  muscles. 

The  Alkali  Metals  as  a  Group.  —  The  alkali  metals  form  a 
natural  group,  the  properties  of  whose  members  exhibit  a  regu- 
lar variation  with  the  atomic  weights.  The  relations  that 
obtain  may  be  seen  from  the  following  table :  — 


ELEMENT 

ATOMIC  WEIGHT 

MELTING  POINT 

SPECIFIC  VOLUME 

Lithium,  Li  

6.94 

180.0° 

11  9 

Sodium,  Na  

2300 

976° 

23  7 

Potassium,  K     .     .     .     . 
Rubidium,  Rb    .... 
Csssium  Cs 

39.10 
85.45 
132  81 

62.5° 

38.5° 
26  5° 

45.2 
56.2 

707 

The  gradation  of  chemical  properties  has  already  been  men- 
tioned. It  is  evident  that  as  the  atomic  weight  increases  the 
chemical  activity  of  the  members  of  the  group  also  increases. 

Spectrum  Analysis.  —  A  little  of  any  sodium  compound  in- 
troduced into  the  Bunsen  flame  colors  it  a  bright  yellow. 
Similarly  potassium  compounds  produce  a  violet  and  lithium 
compounds  a  characteristic  red  color.  These  colored  flames 
may  serve  to  detect  the  presence  of  small  amounts  of  these 
metals.  At  the  high  temperature  of  the  Bunsen  flame,  the 
salts  of  these  and  other  metals  are  decomposed,  and  the  incan- 
descent vapors  emit  light  of  definite  wave  length,  i.e.  of  definite 
color,  in  each  individual  case.  When  sodium  and  potassium 
are  both  present  in  the  flame,  the  latter  appears  yellow,  for  that 
color  completely  masks  the  delicate  violet  of  the  potassium 


366 


OUTLINES   OF   CHEMISTRY 


flame.  The  latter,  however,  can  be  detected  by  viewing  the 
yellow  flame  through  a  blue  glass  or  through  a  layer  of  a  solu- 
tion of  indigo,  in  which  case  the  yellow  light  is  absorbed  and  the 
violet  color  due  to  the  potassium  becomes  apparent.  Robert 
Bunsen  sought  to  work  out  a  system  of  detecting  the  presence 
of  metals  by  noting  the  color  which  they  impart  to  the  flame. 
Later  Bunsen  and  Kirchhoff  (1859)  passed  the  composite  light 
through  a  glass  prism,  and  thus  the  different  colors  could  be 


FIG.  128. 

seen  side  by  side  instead  of  superimposed,  for  light  of  different 
color  possesses  different  refrangibility.  The  instrument  de- 
signed by  Bunsen  and  Kirchhoff  for  producing  and  inspecting 
such  spectra  is  called  a  spectroscope.  Figure  128  shows  a 
simple  form  of  the  instrument  with  the  cover  removed  from  the 
prism.  The  colored  light  from  the  flame  enters  a  narrow  adjust- 
able slit  at  S,  and  a  lens  or  set  of  lenses  in  0  gathers  the  rays 
and  transmits  them  to  the  prism.  After  passing  the  latter,  they 
enter  the  telescope  T,  which  can  be  moved  so  as  to  catch  the  rays. 
Thus,  if  the  flame  is  colored  by  sodium  alone,  the  observer  sees 


THE  ALKALI  METALS 


367 


3l 

g 
a 

A  BC      B        E        f 

•»           ^«   '    '  An          *rt     '     ori 

rA                 tut 
y                   tin 

i       100  —  i  1PQ  l  i  MQ  —  i  IRHi  . 

)           ^0     ,     ^0 

T* 

i      T 

d 

.; 

)n 

w§e^ 

ellow  G 

re 

en 

Blue     In 

ffeo   Vic 

let 

Na 
Li 
K 
Rb 
Cs 
Ca 
Sr 
Ba 
Tl 
In 

1 

i          |                    i                   i 
I          1                    i                   i 

1 

i                                                                      i                        l 

ii 

II                           ii                        I 
ii                           lit                       i 

i         !        !        i        .        I              i 

i     i            i          i          i                    i 
i                  iii                    i                   ! 

i      i 

ill                        i                      i 
I            i                        1                      » 

il                        I                      1 
!             !                          1                        i 

II 

i 

i             i                          i 
I            !                        i                      1 

i 

!        !        !               !              ! 

FIG.  129. 


368  OUTLINES   OF   CHEMISTRY 

simply  one  yellow  line,  the  image  of  the  slit  magnified,  of 
course,  by  the  lenses  of  the  telescope.  If  now  potassium  be 
added  to  the  flame,  the  yellow  line  due  to  the  sodium 
remains,  but  there  appear  in  addition  two  characteristic 
red  lines  to  the  left  of  the  sodium  line,  while  to  the  ex- 
treme right  of  the  latter  a  violet  line  is  found.  The 
two  red  lines  and  the  violet  line  are  due  to  potassium. 
They  always  appear  when  potassium  is  present  and  are 
located  in  the  same  relative  position  to  the  sodium  line 
and  to  one  another.  In  order  to  observe  more  accu- 
rately the  relative  positions  of  these  lines  a  scale  is  re- 
quired. In  the  tube  IF  a  photographic  scale  is  placed 
which  is  illuminated  by  a  candle  or  other  luminous  flame 
as  shown.  Light  from  this  scale  strikes  the  prism  at 
such  an  angle  as  to  be  reflected  through  the  telescope 
to  the  eye.  Thus,  the  scale  and  the  spectrum  are  seen 
together.  Every  incandescent  gas  produces  its  own  char- 
acteristic lines  in  the  spectrum,  and  this  fact  constitutes 
the  basis  of  spectrum  analysis.  Figure  129  exhibits  the 
spectra  of  a  number  of  common  elements,  the  color  of 
the  lines  being  evident  from  the  solar  spectrum  given 
at  the  head  of  the  figure.  Very  minute  quantities  of 
many  substances  can  still  be  detected  by  means  of  the 
spectroscope;  so  one  three-millionth  of  a  milligram  of 
sodium  will  still  show  its  characteristic  yellow  line  in 
the  spectrum. 

When  a  solid  body  is  heated  to  incandescence  in  the 
flame,  the  spectrum  observed  is  not  a  band  spectrum, 
that  is,  it  is  not  made  up  of  a  series  of  definite  lines 
or  bands,  but  consists  of  a  continuous  spectrum  like  that 
produced  when  sunlight  passes  through  the  slit.  It  is 
well  known  that  sunlight  passed  through  a  prism  is 
decomposed,  yielding  a  continuous  spectrum  consisting 
of  the  colors  of  the  rainbow. 

The  light  of  the  sun  thus  yields  a  continuous  spectrum;  this 
is,  however,  crossed  by  numerous  dark  lines  called  the  Fraun- 
hofer  lines.  These  dark  lines  are  produced  by  light  from  the 
incandescent  gases  in  the  sun  passing  through  the  sun's  atmos- 
phere, which  absorbs  the  light  emitted  by  the  gases,  thus  leav- 
ing black  lines  wherever  colored  lines  would  have  been.  The 


THE  ALKALI  METALS 


369 


Fraunhofer  lines,  then,  are  absorption  spectra.  When,  for  exam- 
ple, sodium  light  is  passed  through  the  slit  of  the  spectroscope, 
we  get  its  characteristic  yellow  line  ;  but  if  the  light  is  passed 
through  an  atmosphere  of  sodium  vapor  before  entering  the 
slit,  we  see  a  dark  line  just  where  the  yellow  line  was  before. 
This  is  because  sodium  vapor  absorbs  the  light  produced  by 


800      700 


600 


500 


450 


400 


Fted    Orange  Yellow 

i 

Gnfen 

1 

Blue     1ndi< 

IQ    Violet 

i 

1 

i 
1 

1 

H, 


He 


Vapor 


FIG.  131. 

incandescent  sodium  vapor,  and  thus  we  have  the  dark  line, 
the  so-called  reversed  spectrum.  Now  the  Fraunhofer  lines 
are  similar  lines,  for  the  sun's  atmosphere  contains  the  vapors 
of  the  very  substances  whose  incandescent  gases  near  the  sun 
are  sending  out  light.  The  principle  established  by  Bunsen 
and  Kirchhoff  is  that  the  vapors  of  a  substance  absorb  the  same 
light  which  the  incandescent  gases  of  the  substances  emit.  Thus 
it  has  been  possible  to  analyze  the  sun  and  other  heavenly  bod- 
ies that  emit  light  of  their  own  by  means  of  the  spectroscope. 
The  result  has  been  that  it  has  been  established  that  these  bodies 

2s 


370 


OUTLINES   OF   CHEMISTRY 


contain  practically  the  same  elements  that  are  found  on  the 
earth ;  though  in  some  cases  fewer  and  in  other  cases  more 
lines  have  been  found  in  the  spectra  of  celestial  bodies  than 
correspond  to  known  terrestrial  elements. 

By  means  of  the  spectroscope,  a  number  of  new  elements 
have  been  discovered,  among  which  are  rubidium,  caesium, 
thallium,  indium,  gallium,  helium,  neon,  crypton,  and  xenon. 

Many  metals  whose  salts  cannot  be  vaporized  in  the  Bunsen 
flame  are  heated  in  the  electric  arc.  Spectra  so  obtained  are 
called  spark  spectra.  Again,  gases  are  inspected  spectroscop- 
ically  by  introducing  them  into  a  tube  (Fig.  130),  provided 
with  platinum  or  aluminum  electrodes,  and  then  exhausting 

with  a  vacuum  pump  till  the 
gas  has  a  pressure  of  but  a 
fraction  of  a  millimeter.  When 
such  a  tube  is  connected  with 
an  induction  coil,  the  gas  emits 
light  and  can  thus  be  examined 
with  the  spectroscope  in  the 
usual  way.  Figure  131  gives 
the  spectra  of  a  few  simple 


RED    YELLOW  GREEN    BLUE- 

FIG.  132. 


INPIG.O 


gases. 


AH  colored  solutions  have  their  own  characteristic  absorption 
spectra.  So  if  light  from  a  Welsbach  burner  is  passed  through 
a  colored  solution  and  then  through  the  slit  of  the  spectroscope, 
it  will  appear  that  the  continuous  spectrum  is  crossed  by  black 
absorption  bands  characteristic  of  the  solution.  For  example, 
blood  yields  the  bands  shown  in  Fig.  132,  and  thus  it  is  clear 
that  the  spectroscope  may  be  used  to  detect  the  presence  of 
blood. 

Ammonium  Salts.  —  In  their  general  behavior  these  are  simi- 
lar to  the  salts  of  sodium  and  potassium.  All  ammonium  salts 
are  volatile,  however,  when  heated,  decomposing  into  ammonia 
and  the  acid,  or  into  other  products  of  greater  or  less  complex- 
ity. So  when  ammonium  chloride  NH4C1  is  heated,  it  dissoci- 
ates into  ammonia  and  hydrochloric  acid;  ammonium  nitrite 
NH4NO2  similarly  yields  nitrogen  and  water;  ammonium  ni- 
trate NH4NO3  yields  nitrous  oxide  and  water;  and  ammonium 
oxalate  (NH4)2C2O4  decomposes  into  ammonia,  water,  carbon 
monoxide,  and  carbon  dioxide.  Ammonium  salts  may  be  ob- 


THE  ALKALI  METALS  371 

tained  by  neutralizing  solutions  of  the  acids  with  ammonia  dis- 
solved in  water  and  evaporating  to  dryness.  They  may  fre- 
quently also  be  formed  by  direct  union  of  ammonia  with  the 
acid  in  question.  When  treated  with  an  hydroxide  of  an 
alkali  metal,  ammonium  salts  are  decomposed,  ammonia  being 
liberated. 

It  has  already  been  stated  that  ammonium  chloride  is  ob- 
tained as  a  by-product  in  the  manufacture  of  coal  gas.  Am- 
monium chloride,  or  sal  ammoniac  NH4C1,  crystallizes  in  cubes 
and  dissolves  readily  in  water  with  absorption  of  heat.  At  0°, 
100  parts  of  water  dissolve  28  parts  of  the  salt,  and  at  100°,  73 
parts.  The  salt  has  a  very  sharp,  salty  taste.  Ammonium 
chloride  is  used  in  medicine,  in  making  certain  kinds  of  electric 
batteries,  in  the  dyestuff  industry,  and,  in  general,  as  a  source 
of  ammonia.  Ammonium  bromide  and  ammonium  iodide  are 
deliquescent  salts  that  also  crystallize  in  cubes.  They  slowly 
decompose  on  exposure  to  the  air. 

Ammonium  sulphate  (NH4)2SO4  is  a  very  common  ammonium 
salt,  being  generally  manufactured  by  neutralizing  the  gas 
liquors  with  sulphuric  acid.  It  crystallizes  in  the  rhombic 
system,  and  is  soluble  in  about  1.5  parts  of  cold  water.  The 
salt  is  used  as  a  source  of  ammonia  for  making  other  compounds, 
and  also  as  a  fertilizer.  Over  600,000  tons  of  this  salt  are 
manufactured  annually. 

On  electrolysis  of  a  concentrated  solution  of  ammonium  bisul- 
phate  NH4HSO4,  ammonium  persulphate  (NH4)2S2O8  is  pro- 
duced. The  latter  forms  monoclinic  crystals  that  separate  out 
from  the  solution.  It  serves  as  an  oxidizing  agent. 

Ammonium  sulphide  (NH4)2S  is  obtained  by  passing  hydro- 
gen sulphide  into  a  solution  of  ammonia  in  water,  till  the  latter 
is  half  saturated,  thus  :  — 

2NH3  +  H2S  =  (NH4)2S. 

On  passing  the  hydrogen  sulphide  into  the  solution  till  it 
is  completely  saturated,  ammonium  hydrosulphide  NH4SH 
results :  — 

(NH4)2S  +  H2S  =  2NH4SH. 

The  sulphide  (NH4)2S  may  be  obtained  in  form  of  colorless 
needles  which  readily  give  off  ammonia  and  thus  pass  over  into 
the  hydrosulphide  NH4SH,  which  also  forms  colorless  crystals 


372  OUTLINES   OF  CHEMISTRY 

that  dissociate  into  ammonia  and  hydrogen  sulphide  even  at 
room  temperatures.  At  50°  this  dissociation  is  nearly  complete. 
Almost  invariably  aqueous  solutions  of  ammonium  sulphide  or 
hydrosulphide  are  prepared  for  laboratory  use  as  above  stated. 
These  are  important  in  analytical  chemistry.  Ammonium  sul- 
phide solutions  are  colorless  when  freshly  prepared.  They  have 
a  disagreeable  odor,  due  to  the  fact  that  by  hydrolysis  both 
ammonia  and  hydrogen  sulphide  are  liberated.  On  standing  in 
the  air,  the  solutions  turn  yellow  on  account  of  oxidation.  The 
sulphur  thus  liberated  remains  in  solution,  forming  a  yellow 
polysulphide.  By  dissolving  sulphur  in  ammonium  sulphide 
solutions,  a  series  of  polysulphides  may  be  obtained  which  are 
analogous  to  the  polysulphides  of  the  alkali  metals.  The  solu- 
tions of  these  polysulphides  are  commonly  termed  yellow  ammo- 
nium sulphide.  It  serves  in  analytical  chemistry  for  dissolving 
the  sulphides  of  arsenic,  antimony,  tin,  gold,  and  platinum,  with 
which  it  forms  ammonium  sulpho-salts. 

Ammonium  nitrate  NH4NO3  forms  rhombic  prisms  that  are 
isomorphous  with  potassium  nitrate.  It  melts  at  about  160°, 
and  on  further  heating,  it  decomposes  into  water  and  nitrous 
oxide.  In  water  it  dissolves  readily  with  absorption  of  heat, 
and  it  is  consequently  sometimes  used,  mixed  with  ice,  to  pro- 
duce low  temperatures.  The  salt  is  also  employed  in  explosives 
in  place  of  potassium  nitrate.  Ammonium  nitrite  forms  deli- 
quescent crystals  that  readily  decompose  into  water  and  nitro- 
gen, even  in  aqueous  solutions,  on  being  heated  to  70°. 

Ammonium  carbonate  (NH4)2CO3  +  H2O  is  obtained  in  form 
of  a  crystalline  precipitate  by  passing  carbon  dioxide  into  a 
concentrated  solution  of  commercial  ammonium  carbonate. 
The  latter  consists  of  a  mixture  of  acid  ammonium  carbonate 
NH4HCO3  and  ammonium  carbamate  NH2  •  CO2NH4,  and  is 
obtained  by  heating  either  ammonium  sulphate  or  ammonium 
chloride  with  calcium  carbonate.  The  sublimate  formed  is  a 
hard  white  mass.  The  normal  salt,  (NH^COg,  readily  loses 
ammonia  and  passes  over  into  the  acid  salt,  NH4HCO3,  which 
forms  crystals  that  decompose  into  ammonia,  carbon  dioxide, 
and  water  at  60°.  In  aqueous  solutions,  the  salt  loses  carbon 
dioxide,  thus  forming  the  normal  carbonate. 

Detection  of  Ammonium  Salts.  —  These  salts  are  characterized 
by  their  volatility  and  the  fact  that  ammonia  is  evolved  by 


THE  ALKALI  METALS  373 

treating  them  with  caustic  alkalies.  With  ammonium  chloride, 
platinic  chloride  forms  ammonium  platinic  chloride  (NH4)2PtCl6, 
which,  like  the  analogous  potassium  salt,  is  sparingly  soluble. 
Acid  ammonium  tartrate  C6H6O6 .  NH4  is  precipitated  from 
concentrated  solutions  of  ammonium  salts  by  means  of  tartaric 
acid. 


CHAPTER  XXII 

THE  ALKALINE  EARTH  METALS 

THE  metals  of  the  alkaline  earths  are  calcium,  strontium,  and 
barium.  They  form  another  natural  group  of  closely  related 
elements.  These  metals  never  occur  in  the  free  state  in  nature. 
They  are  harder  than  the  alkali  metals,  have  higher  atomic 
weights,  and  do  not  melt  below  red  heat.  They  act  on  water, 
yielding  hydroxides  and  hydrogen,  though  the  action  is  less 
vigorous  than  in  the  case  of  the  alkali  metals.  The  hydroxides 
formed  are  alkaline  and  rather  sparingly  soluble  in  water,  the 
solubility  increasing  as  the  atomic  weight  of  the  metal  increases. 
On  heating  the  hydroxides,  they  lose  water,  forming  the  oxides, 
which  are  white,  earthy  powders  that  give  an  alkaline  reaction 
with  moist  litmus  paper.  This  dehydration  is  accomplished 
most  readily  in  the  case  of  calcium  hydroxide,  and  least  readily 
in  the  case  of  barium  hydroxide ;  that  is,  the  stability  of  the 
hydroxides  increases  with  the  atomic  weight  of  the  metal. 

In  all  their  compounds  calcium,  strontium,  and  barium  are 

ii 
bivalent.      The   chlorides,   bromides,   and   iodides,  MX2,    the 

ii  ii 

nitrates,  M(NO3)2,  and  the  acetates,  M(C2H3O2)2,  are  readily 

soluble  in  water ;  whereas  the  sulphates,  MSO4,  phosphates, 
M3(PO4)2,  carbonates,  MCO3,  silicates,  MSiO3,  oxalates,  MC2O4, 
and  fluorides,  MF2,  are  sparingly  soluble  in  water.  It  will  be 
recalled  that  the  carbonate  and  phosphate  of  lithium  are  also 
but  slightly  soluble,  and  thus  lithium  in  a  way  represents  a 
transition  between  the  alkali  metals  and  those  of  the  alkaline 
earths.  The  chlorides  of  the  alkaline  earth  metals  are  more 
soluble  than  the  nitrates.  The  solubility  of  the  sulphates 
decreases  as  the  atomic  weight  increases.  The  insoluble  car- 
bonates, sulphates,  phosphates,  and  silicates  are  specially  charac- 
teristic of  this  group.  They  are  of  importance  in  nature,  in  the 
arts  and  industries,  and  in  chemical  analysis.  The  acid  car- 
bonates or  bicarbonates  are  much  more  soluble  than  the 
carbonates. 

374 


THE  ALKALINE  EARTH  METALS 


375 


Occurrence,  Preparation,  and  Properties  of  Calcium.  —  This 
metal  occurs  very  widely  distributed  and  often  in  enormous 
masses  as  carbonate  in  form  of  marble,  chalk,  or  limestone.  It  is 
further  found  as  sulphate  in  form  of  gypsum  and  anhydrite,  as 
phosphate,  as  fluoride  or  fluorspar,  and  as  an  essential  constit- 
uent of  many  silicates.  Nearly  all  natural  waters  contain 
calcium  sulphate  and  bicarbonate.  The  bones  and  teeth  of 
animals  consist  mainly  of  calcium  phosphate  together  with  some 
carbonate  and  small  amounts  of  fluoride.  In  eggshells,  coral, 
and  the  shells  of  molluscs  and  various  crustaceans  the  carbon- 
ate of  calcium  predominates.  Calcium  salts  are  also  found  in 
plants.  These  are  the  sulphate,  oxalate,  phosphate,  and  car- 
bonate, as  well  as  salts  of  various  complex  organic  acids. 
Similarly  calcium  compounds  are  distributed  throughout  the 
bodies  of  animals  in  small  quantities. 

Metallic  calcium  may  be  obtained  by  heating  calcium  iodide 
with  sodium,  or  by  heating  an  excess  of  calcium  oxide  with 
carbon  or  calcium  carbide  in  the  A 
electric  furnace.  The  best  way  to 
prepare  the  metal  is  by  electroly- 
sis of  the  molten  chloride.  The 
chloride  is  placed  in  a  carbon  con- 
tainer (Fig.  133),  the  walls  of 
which  serve  as  anode  (see  elec- 
trolysis). The  cathode  is  of  iron 
or  copper.  After  the  electrolysis 
has  started,  the  heat  developed  by 
the  current  is  sufficient  to  keep 
the  salt  in  molten  condition.  The 
metal  separates  out  at  the  cathode, 
and,  being  light,  rises  to  the  top 
of  the  molten  chloride.  By  slowly 
raising  the  cathode  as  the  elec- 
trolysis proceeds,  a  rough  stick  of 
calcium  is  obtained,  for  the  metal 
adheres  to  the  electrode.  Metallic 
calcium  may  now  be  purchased  at  less  than  a  dollar  a  pound. 

Calcium  is  a  silver- white  metal  of  specific  gravity  1.85. 
It  crystallizes  in  the  hexagonal  system  and  is  fairly  hard, 
tough,  and  malleable.  It  may  readily  be  worked  in  a  lathe, 


Vteter 


FIG.  133. 


376  OUTLINES  OF  CHEMISTRY 

It  decomposes  water,  and  consequently  is  kept  under  petroleum, 
or  more  frequently  simply  in  air-tight  containers  of  glass  or 
tinned  iron.  Calcium  melts  at  about  760°,  and  at  that  tem- 
perature catches  fire  in  the  air,  burning  to  the  oxide,  CaO,  and 
the  nitride,  Ca3N2.  The  latter  is  a  yellow  powder  which  is  de- 
composed by  water,  yielding  the  hydroxide  and  ammonia. 
Calcium  was  formerly  described  as  a  yellow  metal.  The  yellow 
color  was  due  to  the  presence  of  calcium  nitride  as  an  impurity. 
With  hydrogen,  calcium  readily  forms  calcium  hydride  CaH2, 
a  white  powder  that  acts  more  vigorously  on  water  than  the 
metal  itself.  In  general,  calcium  is  very  active  chemically,  for 
it  unites  with  all  the  elements  except  those  of  the  argon  group. 
Calcium  Carbonate  CaCO3  is  the  most  abundant  of  all  the  cal- 
cium compounds.  In  an  impure  form,  as  limestone,  it  forms 
mountains  and  strata  of  vast  extent  and  great  thickness.  Dolo- 
mite is  essentially  a  calcium,  magnesium  limestone  of  the  com- 
position MgCO3  •  CaCO3.  It  generally  contains  silica,  iron, 
alumina,  and  other  impurities.  Marl  consists  of  limestone 
mixed  with  clay.  Chalk  is  fairly  pure  calcium  carbonate.  In 
a  crystalline  state  calcium  carbonate  occurs  as  marble  in  many 
localities.  In  pure  form  it  occurs  as  calcite,  particularly  as  Ice- 
land spar  in  Iceland.  The  stalactites  and  stalagmites  found  in 
many  caves  consist  of  calcium  carbonate.  Calcium  carbonate 
crystallizes  in  the  hexagonal  system  as  calcite,  commonly  form- 
ing rhombohedra  (Figs.  65  and  66)  or  scalenohedra  (Figs.  64 
and  67),  and  also  in  the  orthorhombic  system  as  aragonite. 
The  hexagonal  form  is  the  stable  one  at  ordinary  temperatures ; 
at  higher  temperatures  the  rhombic  form  is  stabler.  So  from  a 
cold  solution  calcium  carbonate  deposits  in  hexagonal,  and  from 
a  hot  solution  in  rhombic,  crystals.  Calcium  carbonate  is  prac- 
tically insoluble  in  water;  but  in  water  charged  with  carbon 
dioxide  it  dissolves  fairly  readily.  The  solution  contains  cal- 
cium bicarbonate  Ca(HCO3)2,  in  all  probability.  On  boiling 
such  solutions,  carbon  dioxide  is  expelled  and  the  normal 
carbonate  is  precipitated.  Waters  containing  calcium  salts  in 
solution  are  said  to  be  hard.  The  hardness  that  can  be  dis- 
pelled by  boiling,  as  just  mentioned,  is  termed  temporary  hard- 
ness, as  compared  with  permanent  hardness  which  is  produced 
by  the  presence  of  calcium  salts  other  than  the  bicarbonate,  and 
consequently  persists  even  on  boiling. 


THE  ALKALINE  EARTH  METALS 


377 


Limestone  and  marble  are  used  as  building  stones,  in  glass 
manufacture,  in  the  reduction  of  iron  ores,  and  in  making  lime, 
Portland  cement,  sodium  carbonate,  and  many  other  products. 
Much  calcium  carbonate  is  also  used  as  chalk  and  whiting. 
Mixed  with  linseed  oil,  calcium  carbonate  forms  putty. 

Calcium  Oxide  CaO,  lime,  is  formed  by  heating  calcium  car- 
bonate above  600°.  To  obtain  the  pure  oxide,  pure  calcium 
carbonate  is  strongly  ignited  ;  whereas  on  a  commercial  scale, 
lime  is  prepared  by  heating  limestone  in  limekilns  (Fig.  134). 
Lime  is  a  white,  amorphous, 
porous  solid.  It  may  be  melted 
in  the  electric  furnace.  Lime 
unites  with  water  with  evolu- 
tion of  heat,  forming  a  powder, 
slaked  lime  or  calcium  hydroxide 
Ca(OH)2.  The  unslaked  oxide, 
CaO,  is  called  quicklime  or 
caustic  lime.  Calcium  hydrox- 
ide is  somewhat  soluble  in 
water.  At  15°,  100  parts  of 
water  dissolve  about  0.14  part 
of  the  hydroxide,  while  at  100° 
but  half  of  that  amount  dis- 
solves. The  solution  is  alka- 
line and  is  known  as  limewater, 
when  clear.  When  it  contains  excess  of  hydroxide  in  suspen- 
sion, it  is  termed  milk  of  lime.  When  exposed  to  the  air, 
quicklime  gradually  absorbs  moisture  and  carbon  dioxide,  and 
crumbles,  being  converted  into  calcium  carbonate  or  air-slaked 
lime. 

Besides  being  used  in  making  mortar  for  building  purposes, 
lime  is  generally  employed  in  chemical  industries  when  a  cheap 
base  is  required.  Large  amounts  are  used  in  purifying  coal  gas 
and  sugar,  in  removing  hair  from  hides,  in  making  bleaching 
powder,  sodium  and  potassium  hydroxides,  glass,  and  oxalic, 
tartaric,  and  citric  acids.  Lime  is  also  used  as  a  disinfectant, 
and  limewater  is  frequently  employed  in  medicine. 

Mortar  consists  of  a  pasty  mass  obtained  by  mixing  sand, 
slaked  lime,  and  water.  After  it  has  been  applied,  water  dries 
out  gradually  and  carbon  dioxide  is  absorbed  from  the  air,  thus 


FIG.  134. 


378  OUTLINES  OF  CHEMISTRY 

forming  hard  calcium  carbonate.  This  process  is  called  the 
setting  of  the  mortar.  It  will  not  take  place  while  the  mortar 
remains  wet,  and  will  commonly  require  a  rather  long  time  for 
its  completion;  for  after  the  outer  layer  has  become  trans- 
formed to  carbonate,  the  deeper  layers  are  partially  protected 
from  the  air  and  so  are  altered  but  slowly.  When  thoroughly 
hardened,  the  sand  grains  are  firmly  fixed  in  the  matrix  of 
crystalline  calcium  carbonate,  which  adheres  well  to  the  brick 
or  stone  used.  Lime  mortar  is  unfit  for  use  in  places  that  are 
always  wet,  for  it  hardens  only  when  dry.  After  hundreds  of 
years,  some  calcium  silicate  is  formed  by  interaction  of  the  lime 
with  the  sand  grains.  Lime  mortar  has  been  in  common  use 
for  a  very  long  time.  It  was  quite  generally  employed  by  the 
Romans  in  their  buildings. 

Lime  made  from  magnesian  limestone,  dolomite,  consists  of 
CaO  and  MgO.  It  slakes  very  slowly  with  cold  water,  owing 
to  the  presence  of  the  magnesia.  In  cold  weather,  hot  water 
is  commonly  employed  in  slaking  this  lime,  which,  however, 
makes  very  good  mortar  and  hence  is  used  in  many  localities. 

Cement.  —  When  limestone  containing  silicates  of  aluminum 
is  heated  in  a  kiln  and  the  product  ground  to  a  powder,  the 
latter  forms  a  so-called  hydraulic  cement,  for  it  will  unite  with 
water  and  form  a  hard,  insoluble  mass.  The  hardening  process 
takes  place  uniformly,  and  relatively  quickly  throughout  the 
mass  even  under  water.  The  cement  is  very  valuable,  for  it 
can  be  used  in  damp  as  well  as  in  dry  places.  In  some  locali- 
ties, as  near  Milwaukee  and  Louisville,  limestone  containing 
a  suitable  amount  of  aluminum  silicates  for  producing  cement 
is  found  and  made  into  cement.  Such  cements  are  called 
natural  cements.  They  generally  contain  notable  quantities  of 
magnesia  and  other  ingredients.  The  following  table  gives  the 
composition  of  Louisville  cement  in  per  cent :  — 

SiO2.     ...     .  21.1  CaSO4    ....  6.8 

A1208J  K20   | 

Fe208|  '     '     •     '  7'5  Na2OJ    '     '     '     '  °'8 

CaO  .     .     .     .     .  44.4  CO2 11.2 

MgO      ....  7.0  H2O 1.2 

Portland  cement  is  made  by  artificially  mixing  materials  con- 
taining silica,  lime,  and  alumina  in  proper  proportions  and  then 


THE  ALKALINE  EARTH  METALS  379 

firing  the  mixture  in  a  kiln.  The  hard  mass  or  clinker  thus 
obtained  is  ground  to  a  fine  powder  and  constitutes  the  so-called 
Portland  cement.  In  practice,  clay  rich  in  silica,  and  calcium 
carbonate,  are  used  in  making  this  cement.  Often  marl  or  the 
slag  from  blast  furnaces  is  employed.  In  all  cases,  it  is  neces- 
sary to  determine  the  composition  of  the  materials  used  by 
chemical  analysis,  so  that  they  may  be  mixed  in  the  proper 
proportions  of  silica,  alumina,  and  lime.  The  following  table 
gives  the  composition  of  Portland  cement  in  per  cent.  The 
last  column  indicates  the  result  of  an  analysis  of  a  typical 
American  Portland  cement:  — 

COMPOSITION  OF  PORTLAND   CEMENT  IN  PER  CENT 


Si02    .     .     . 
CaO     .     .-   . 

.     .     .     .     20  to  25 
.     .     .     .     58  to  65 

22.5 
62.9 

ALOo  . 

.     .     .     .       5  to  10 

6.4 

Mg-0   . 

traces  to    5 

3.1 

.     .     ,     .       2  to    5 

3.3 

KO  J-  IVa   O 

H20  +  C02  . 

SOo 

.    traces  to    2 
traces  to    3 

0.7 
1.0 

99.9 

No  definite  chemical  formula  can  be  ascribed  to  Portland  cement, 
the  composition  of  excellent  cements  varying  within  the  limits 
indicated  in  the  first  column  above.  Experience  has  shown  that 
the  ratio  of  the  amount  of  lime  to  that  of  silica  plus  alumina 
and  ferric  oxide  should  fall  between  1.8  and  2.2,  being  about 
2  on  the  average.  That  is, 

_  CaO  _ 

SiO2  +  A12O3  +  Fe2O3 

should  be  greater  than  1.8  and  less  than  2.2;  for  the  cement 
whose  analysis  is  given  in  the  table  above,  the  ratio  is 


=1.95. 


22.5+6.4+3.3 

The  chemical  changes  involved  in  the  making  and  hardening 
of  cement  have  been  the  subject  of  much  investigation  and 
discussion.  The  clinker  probably  contains  mainly  tricalcium 


380  OUTLINES   OF   CHEMISTRY 

silicate  Ca3SiO5  and  calcium  aluminate  Ca3Al2O6.  These  when 
ground  fine  and  then  treated  with  water  probably  suffer  decom- 
position thus :  — 

2  Ca3Si05  +  9  H20  =  4  Ca(OH)2  +  (CaSiO3)2  -  5  H2O,  and 
Ca3Al206  +  Ca(OH)2  +  11  H2O  =  Ca4Al2O7  . 12  H2O, 

so  that  the  hardened  cement  contains  the  hydrated  calcium 
silicate  (CaSiO3)2  •  5  H2O  and  the  hydrated  basic  calcium  alu- 
minate Ca4Al2O7  •  12  H2O.  The  hardening  of  the  cement  is 
supposed  to  be  due  mainly  to  the  formation  of  the  former 
compound.  Cement  powder  has  a  greenish  gray  color.  Its 
specific  gravity  is  3.1  to  3.2.  The  hardened  mass  has  a  drab 
color,  resembling  the  rock  found  at  Portland,  England,  whence 
the  name  of  the  cement.  Mixed  with  crushed  stone  and  water 
in  proper  proportions,  Portland  cement  hardens  into  a  mass 
called  concrete,  which  wears  as  well  as  excellent  stone.  Often 
concrete  is  strengthened  by  imbedding  in  it  rods  of  iron  or 
steel.  The  product  is  called  reenforced  concrete.  Plain  and 
reenforced  concrete  are  much  used  in  modern  structures,  and 
the  production  of  cement  has  greatly  increased  in  recent  years. 
In  1910,  14,779,177  tons  of  cement  were  made  in  the  United 
States,  which  is  more  than  twice  the  amount  produced  in  1904. 
Calcium  Sulphate  CaSO4  occurs  in  large  quantities  in  nature 
as  gypsum  CaSO4  •  2  H2O,  which  forms  monoclinic  crystals 
(Fig.  73)  called  selenite,  and  also  as  anhydrite  CaSO4  in  rhom- 
bic forms.  Alabaster  is  a  granular,  crystalline  form  of  gypsum. 
Calcium  sulphate  occurs  in  soils  and  natural  waters.  At  0°, 
100  parts  of  water  dissolve  0.19  part  of  calcium  sulphate;  at 
35°,  0.21  part.  In  nitric  or  hydrochloric  acids  and  in  many 
salt  solutions  calcium  sulphate  dissolves  much  more  copiously. 
At  about  110°,  gypsum  loses  water,  forming  (CaSO4)2  •  H2O, 
which  is  a  white  powder  known  as  plaster  of  Paris.  When 
this  is  mixed  with  water,  a  paste  may  be  obtained  which  soon 
hardens.  This  hardening  or  "  setting  "  is  due  to  the  fact  that 
water  is  again  taken  up,  a  crystalline,  coherent  mass  of  gypsum 
being  formed,  thus  :  — 

(CaS04)2  •  H20  +  3  H20  =  2 (CaSO4  -  2  H2O). 

Plaster  of  Paris  is  much  used  for  making  casts,  surgical  band- 
ages, and  stucco. 


THE  ALKALINE  EARTH  METALS  381 

At  about  200°,  gypsum  loses  all  of  its  water,  and  is  then  said 
to  be  dead  burned,  for  in  this  condition  it  unites  with  water 
but  slowly  and  without  hardening.  Gypsum  is  often  used  as 
a  fertilizer,  land  plaster.  Its  action  probably  depends  on  the 
fact  that  it  reacts  with  the  ammonium  carbonate  in  soils,  form- 
ing ammonium  sulphate,  which,  being  practically  non-volatile, 
remains  in  the  soil  and  is  utilized  by  plants. 

Calcium  Sulphite  CaSO3  is  formed  by  passing  sulphur  dioxide 
into  calcium  hydroxide.  The  salt  crystallizes  in  prisms  of  the 
composition  CaSO3  •  2  H2O,  which  are  slightly  soluble  in  water, 
1  in  800.  In  aqueous  solutions  of  sulphur  dioxide,  the  salt  dis- 
solves more  copiously,  and  such  solutions  are  used  in  paper 
mills  in  preparing  wood  pulp. 

Calcium  Sulphide  CaS  is  obtained  by  heating  calcium  sul- 
phate with  charcoal :  — 

CaSO4  +  4  C  =  CaS  +  4  CO. 
With  water  the  sulphide  reacts  thus  :  - — 

2  CaS  +  2  H2O  =  Ca(OH)2  -f  Ca(SH)2, 

the  latter  compound  being  soluble  in  water.  Calcium  sulphide 
is  used  in  making  luminous  match  safes,  clock  faces,  etc.,  for, 
after  exposure  to  sunlight,  it  emits  a  faint  light  which  is  visible 
in  the  dark.  Barium  sulphide  BaS  and  strontium  sulphide  SrS 
serve  similarly  for  making  so-called  luminous  paint. 

Calcium  Fluoride  CaF2  crystallizes  in  cubes,  and  is  found  in 
nature  as  fluorite.  It  is  insoluble  in  water.  It  serves  as  a  flux, 
and  is  used  in  making  hydrofluoric  acid  and  other  fluorine  com- 
pounds. 

Calcium  Chloride  CaCl2  occurs  at  Stassfurt  in  tachhydrite 
CaCl2-  MgCl2- 12  H2O.  It  is  obtained  as  a  by-product  in  the 
Solvay  soda  process  and  in  the  production  of  ammonia  by 
the  action  of  lime  on  ammonium  chloride.  It  is  also  made 
by  the  action  of  hydrochloric  acid  on  calcium  carbonate  or 
hydroxide.  From  solutions  it  crystallizes  in  hexagonal  prisms, 
CaCl2 .  6  H2O.  These  melt  in  their  crystal  water  at  29°,  and 
become  a  porous,  anhydrous  mass  at  200°.  The  anhydrous  salt 
melts  at  719°.  It  is  deliquescent,  and  is  much  used  as  a  dry- 
ing agent  in  chemical  laboratories.  With  ice  and  the  hydrate, 
CaCl2  •  6  H2O,  temperatures  as  low  as  —  50°  may  be  produced. 


382  OUTLINES   OF  CHEMISTRY 

The  anhydrous  salt  dissolves  in  water  with  liberation  of  heat. 
With  alcohol,  and  also  with  ammonia,  calcium  chloride  forms 
addition  products,  so  that  these  substances  must  be  dried  with 
other  agents  like  the  oxide  of  calcium  or  barium.  Calcium  bro- 
mide CaBr2  and  calcium  iodide  CaI2  are  even  more  deliquescent 
than  calcium  chloride. 

Bleaching  Powder  or  chloride  of  lime  is  made  in  large  quanti- 
ties by  passing  chlorine  into  calcium  hydroxide,  i.e.  slaked 
lime.  The  composition  of  the  compound  is  expressed  by  the 
formula  Ca(OCl)Cl,  as  explained  under  chlorine,  where  the 
reactions  involved  in  its  preparation  and  use  are  also  de- 
scribed. Bleaching  powder  is  slightly  yellowish  in  color.  It 
absorbs  carbon  dioxide  and  moisture  from  the  air.  Thus  hypo- 
chlorous  acid  is  formed,  to  which  the  odor  of  bleaching  powder 
is  due.  Enormous  quantities  of  bleach,  as  it  is  also  called,  are 
used  in  paper  making  and  irr  the  manufacture  of  cotton  and 
linen  goods. 

Calcium  Phosphate  Ca3  (PO4)2  is  found  in  nature  as  already 
stated,  in  apatite,  and  in  the  bones  of  animals.  By  treating  a 
solution  of  calcium  chloride  with  sodium  ammonium  phosphate, 
calcium  phosphate  is  precipitated,  thus  :  — 

2  Na2NH4PO4  +  3  CaCl2  =  4  NaCl  +  2  NH4C1  +  Ca3(PO4)2. 

Calcium  phosphate  is  practically  insoluble  in  water,  but  in 
acids  it  dissolves  readily,  also  in  many  solutions  of  salts,  like 
chlorides  or  nitrates  of  the  alkalies.  As  the  latter  are  present 
in  soils,  calcium  phosphate  is  dissolved  by  their  solutions  and 
hence  made  available  to  plants.  Calcium  phosphate  is  necessary 
to  plant  life,  and  also  to  the  life  of  animals,  into  whose  bones  it 
enters  as  a  prime  constituent.  So-called  superphosphate  of  lime, 
a  fertilizer  of  great  value,  consists  of  calcium  sulphate  and 
primary  calcium  phosphate  produced  by  the  action  of  sulphuric 
acid  on  calcium  phosphate,  thus  :  — 

Ca3(PO4)2  +  2  H2SO4  =  CaH4(PO4)2  +  2  CaSO4. 

The  primary  calcium  phosphate  CaH4(PO4)2  is  soluble  in 
water  and  hence  is  "readily  available  to  plants. 

In  1910  the  United  States  produced  2,654,988  tons  of  phos- 
phate rock,  most  of  which  was  used  as  fertilizer.  The  most 
important  beds  of  this  rock  are  in  South  Carolina  and  Florida. 
As  this  material  is  of  prime  importance  in  maintaining  the  fer- 


THE  ALKALINE  EARTH  METALS  383 

tility  of  the  soil,  its  exportation  has  recently  been  forbidden 
by  law. 

Calcium  Carbide  CaC2  is  made  by  heating  lime  or  calcium 
carbonate  with  carbon  in  the  electric  furnace,  thus  :  — 

CaO+3C  =  CaC2  +  CO. 

Pure  calcium  carbide  is  white,  but  the  commercial  article  is 
dark  in  color,  owing  to  the  presence  of  impurities.  The  sub- 
stance yields  acetylene  when  treated  with  water,  as  already 
stated,  and  hence  large  quantities  of  it  are  manufactured 
annually. 

Calcium  Phosphide  Ca2P2  is  formed  by  heating  lime  and 
phosphorus  together,  thus  :  — 

14  CaO  +  14  P  =  2  Ca2P2O7  +  5  Ca2P2. 

The  product  is  a  brown  solid,  which  on  treatment  with  water 
yields  calcium  hydroxide  and  phosphine. 

Calcium  Cyanamide  CaN  •  CN  is  formed  by  passing  nitrogen 
over  calcium  carbide  heated  to  white  heat  in  an  electric  furnace, 

thus : — 

CaC2  +  N2  =  CaN  •  CN  +  C. 

Calcium  cyanamide  is  used  as  a  fertilizer,  for  in  the  soil  its 
nitrogen  gradually  is  converted  into  ammonia  and  nitrates, 
owing  to  the  action  of  water  and  oxygen  from  the  air. 

Calcium  Silicide  CaSi2  is  produced  by  heating  lime  with 
silicon  in  the  electric  furnace.  It  forms  hexagonal  crystals 
that  react  but  slowly  with  water.  Dilute  acids  decompose 
the  silicide  readily. 

Calcium  Silicate  CaSiO8  is  occasionally  found  in  nature  in 
monoclinic  crystals  as  wollastonite.  It  occurs  very  frequently 
in  complex  silicates  like  mica,  feldspar,  garnet,  hornblende,  and 
many  ethers.  Calcium  silicate  may  be  formed  by  heating 
together  silica  and  lime  or  calcium  carbonate,  also  by  adding 
sodium  silicate  to  a  solution  of  a  calcium  salt,  thus  :  — 

CaCl2  +  Na2SiO3  =  CaSiO3  +  2  NaCl. 

Calcium  silicate  is  practically  insoluble  in  water  ;  but  in  hydro- 
chloric acid  it  dissolves,  being  decomposed  into  calcium  chloride 
and  silicic  acid. 

Glass  is  a  mixture  of  the  silicates  of  sodium  and  calcium. 
Sodium  silicate  is  transparent  and  soluble  in  water,  while 


384 


OUTLINES   OF   CHEMISTRY 


calcium  silicate  is  neither  transparent  nor  soluble  in  water. 
When  sodium  silicate  is  fused  together  with  calcium  silicate 
in  proper  proportion,  a  liquid  results  which  on  cooling  forms 
ordinary  glass.  This  is  practically  insoluble  in  both  water 
and  acids,  except  hydrofluoric  acid.  Ordinary  glass  varies 
somewhat  in  composition.  It  is  a  mixture  of  silicates  of 
sodium  and  calcium  which  approximates  the  composition 
Na2O  •  3  SiO2  +  CaO  •  3  SiO2.  This  soda-lime  glass,  as  it  is  also 
called,  is  used  for  windows  and  various  ordinary  vessels.  It 
melts  readily  and  is  consequently  relatively  easy  to  work  into 
desired  forms.  In  making  glass,  finely  ground  and  intimately 

mixed  quartz  sand 
SiO2,  calcium  car- 
bonate CaCO3,  and 
soda  Na2CO3  are 
melted  together 
in  large  pots  of 
special  fire  clay. 
These  pots  are 
about  4  feet  high 
and  4  feet  in  di- 
ameter ;  Fig.  135 
shows  an  open  and  a  closed  form.  During  this  process  carbon 
dioxide  escapes.  The  temperature  is  raised  after  the  mass  has 
melted,  to  expel  all  gases.  The  scum  and  various  impurities 
that  have  gathered  on  top  of  the  molten  material  are  removed 
mechanically.  In  making  plate  glass,  the  molten,  viscous 
material  is  poured  upon  large  iron  or  bronze  plates  and  rolled 
to  the  desired  thickness  with  hot  iron  rollers.  The  plates  are 
afterwards  ground  and  polished.  For  ordinary  window  glass, 
the  glass  is  first  blown  into  cylindrical  forms  which  are  cut 
open  lengthwise  and  then  flattened  out.  Heavy  glass  dishes 
and  many  other  similar  articles  are  made  by  pressing  the  plastic 
glass  into  suitable  molds.  Bottles  are  made  by  taking  glass  on 
the  end  of  an  iron  pipe  and  blowing  it  into  molds.  All  flasks, 
retorts,  beakers,  and  many  other  thin  glass  dishes  are  blown 
into  suitable  molds.  Glass  tubing  is  made  by  blowing  a  small 
bulb  on  the  end  of  an  iron  blowpipe,  and  then  attaching  the 
bulb  at  its  lower  end  and  drawing  the  plastic  mass  out  into 
the  form  of  a  tube.  All  glass  articles  must  be  annealed.  This 


FIG.  135. 


THE   ALKALINE   EARTH  METALS  385 

process  consists  of  cooling  the  glass  gradually  in  suitable  fur- 
naces. Glass  suddenly  cooled  is  under  internal  strain,  and  a 
slight  scratch  will  cause  it  to  break  ;  thus  the  "  Prince  Rupert's 
drop  "  of  suddenly  chilled  glass  is  shattered  by  nipping  off  its 
tail.  The  strength  and  transparency  of  glass  depends  upon  its 
amorphous  nature,  for  when  glass  crystallizes  it  becomes  opaque 
and  brittle.  Since  the  composition  of  glass  may  be  varied  to  a 
considerable  extent  at- will,  it  is  of  the  nature  of  a  solution;  and 
indeed  it  is  sometimes  called  a  solid  solution,  resulting  from 
the  supercooling  of  the  liquid  mass.  Since  sodium  silicate  is 
soluble  in  water,  and  calcium  silicate  is  readily  decomposed  by 
hydrochloric  acid,  and  on  the  other  hand  glass  is  practically 
not  attacked  by  water  or  hydrochloric  acid,  it  is  clear  that  in 
glass  the  silicates  of  sodium  and  calcium  are  chemically  com- 
bined with  each  other,  in  spite  of  the  fact  that  their  relative 
proportions  may  gradually  be  varied  to  some  extent. 

Soda-lime  glass  is  unsuitable  for  laboratory  utensils,  for  it  is 
too  easily  attacked  by  chemical  reagents,  especially  by  alkalies. 
For  this  reason  potash-lime  glass,  in  the  manufacture  of  which 
potassium  carbonate  is  used  instead  of  sodium  carbonate,  is 
commonly  employed  in  making  finer  glassware,  for  it  is  much 
harder  and  less  readily  acted  upon  by  reagents.  This  potash 
glass  is  also  called  Bohemian  glass  or  hard  glass ;  it  has  a  much 
higher  melting  point  than  ordinary  soda  glass,  which  is  fre- 
quently called  soft  glass.  Crown  glass  is  a  potash-lime  glass. 
Flint  glass  is  produced  by  melting  together  silica,  potash,  and 
lead  oxide.  It  has  a  high  specific  gravity  and  also  a  high  index 
of  refraction,  hence  it  is  used  in  making  lenses,  prisms,  and 
other  optical  instruments  as  well  as  ornaments.  Cut  glass  is 
made  by  grinding  and  polishing  flint  glass.  Jena  glass,  which 
is  now  much  used  for  utensils,  contains  boric  anhydride.  Colored 
glass  is  made  by  adding  small  quantities  of  various  metallic 
oxides  to  molten  glass.  In  this  way  colored  silicates  are 
formed.  Thus  cobalt  oxide  colors  glass  blue;  chromic  oxide 
produces  a  bright  green  shade ;  manganese  dioxide  yields  a 
violet  color ;  cuprous  oxide  or  gold  produces  a  red  shade  owing 
to  the  very  finely  suspended  particles  of  these  substances  in  the 
glass.  The  ordinary  green  bottle  glass  is  colored  by  ferrous 
oxide,  while  the  brown  or  brownish  yellow  glass  is  colored  by 
ferric  oxide.  Window  glass  also  contains  small  amounts  of 
2c 


386 


OUTLINES  OF  CHEMISTRY 


iron,  which  produces  the  green  color  observed  when  viewing 
the  glass  on  edge.  Black  or  very  dark  glass  is  produced  by 
large  amounts  of  cobalt,  iron,  and  other  oxides.  Bottle  glass 
is  commonly  made  from  impure,  cheap  materials  that  contain 
relatively  large  amounts  of  iron  and  other  oxides  that  color  the 
glass.  Enamel  glass  is  a  lead  glass  to  which  oxide  of  tin  has 
been  added,  while  milk  glass  contains  calcium  phosphate  in 
suspension,  which  produces  the  characteristic  white,  opaque 
appearance. 

The  following  table  gives  the  approximate  percentage  com- 
position of  a  few  typical  kinds  of  glass. 


SiO2 

Na20 

K20 

CaO 

PbO 

t 

Quartz  glass     .... 

100 

_ 

_ 

Common  glass      .     .     . 

75.5 

12.9 

— 

11.6 

— 

Bohemian  or  crown  glass 

70.8 

— 

18.3 

10.9 

— 

Crystal  or  flint  glass 

53.5 

— 

13.8 

— 

32.7 

The  art  of  making  glass  is  very  old.  It  is  supposed  to  have 
originated  with  the  Phoenicians.  The  Egyptians  practiced  it 
long  before  the  Christian  era.  From  the  thirteenth  to  the 
seventeenth  century  Venice  was  noted  for  its  glass  manufac- 
tures. The  general  introduction  of  the  use  of  window  glass 
dates  from  about  the  sixteenth  century.  Bohemian  glass  was 
placed  in  the  market  as  early  as  the  fifteenth  century,  though 
glass  was  manufactured  in  Germany  and  Bohemia  even  in  the 
tenth  century. 

Occurrence,  Preparation,  and  Properties  of  Strontium.  —  Stron- 
tium occurs  in  nature  as  celestite  SrSCX  and  strontianite  SrCOo. 

4  O 

It  is  of  much  rarer  occurrence  than  either  calcium  or  barium. 
The  metal,  which  has  properties  similar  to  those  of  calcium, 
may  be  obtained  by  electrolysis  of  the  molten  chloride,  SrCl2. 
Metallic  strontium  acts  on  water,  liberating  hydrogen  and 
forming  strontium  hydroxide.  In  the  air,  strontium  is  slowly 
oxidized.  The  metal  is  silver  white  and  plastic  like  lead.  It 
melts  at  about  800°  and  volatilizes  readily  at  950°. 

Strontium  Compounds.  —  These  are  in  general  very  similar 
to  the  compounds  of  calcium.  They  are  prepared  by  treating 
the  carbonate  with  acids.  The  native  sulphate  serves  as  the 


THE   ALKALINE   EARTH   METALS  387 

chief  source  of  strontium  salts.  It  is  first  reduced  to  sulphide 
by  heating  with  charcoal  and  then  treated  with  acids.  Stron- 
tium salts  impart  a  brilliant  red  color  to  the  Bunsen  flame. 
The  spectrum  contains  a  blue  line,  an  orange  line,  and  six 
bright  lines  in  the  red.  Strontium  carbonate  is  not  as  readily 
decomposed  into  strontium  oxide  SrO  and  carbon  dioxide  as 
calcium  carbonate.  With  water,  the  oxide  readily  forms  stron- 
tium hydroxide  Sr(OH)2,  which  is  more  soluble  than  calcium 
hydroxide.  With  sugar,  strontium  hydroxide  forms  insoluble 
compounds,  hence  it  is  used  in  extracting  sugar  from  molasses 
that  will  no  longer  yield  crystals  of  sugar.  Strontium  dioxide 
SrO2  is  also  known. 

Strontium  chloride  SrCl2  is  hygroscopic,  but  it  does  not  pos- 
sess this  property  to  as  marked  a  degree  as  calcium  chloride. 
The  salt  is  isomorphous  with  calcium  chloride,  forming  hexag- 
onal prisms  of  the  composition  SrCl2  +  6  H2O. 

Strontium  sulphate  SrSO4  is  precipitated  from  solutions  of 
strontium  salts  by  adding  a  solution  of  a  soluble  sulphate.  It 
dissolves  in  about  7000  parts  of  water.  In  alcoholic  solutions 
it  is  much  less  soluble.  On  boiling  strontium  sulphate  with 
sodium  or  potassium  carbonate,  the  salt  is  transformed  into  the 
carbonate :  — 

SrSO4  +  Na2CO3  =  Na2SO4  +  SrCO3. 

Under  like  treatment,  barium  carbonate  remains  practically 
unchanged. 

Strontium  nitrate  Sr(NO3)2  forms  octahedra  or  cubes  when 
crystallized  from  concentrated  solutions.  At  low  temperatures 
the  hydrate,  Sr(NO3)2  •  4  H2O,  separates  from  solutions  in  mono- 
clinic  crystals.  These  effloresce  on  exposure  to  the  air.  Though 
readily  soluble  in  water,  strontium  nitrate  is  insoluble  in  alcohol. 
When  heated,  the  nitrate  is  decomposed,  yielding  the  oxide. 
Strontium  nitrate  is  much  used  in  fireworks  and  red  Bengal  lights. 
The  latter  usually  consist  of  a  mixture  of  about  50  parts  stron- 
tium nitrate,  30  parts  shellac,  and  20  parts  potassium  chlorate. 

Besides  being  used  in  the  sugar  industry,  strontium  salts  are 
sometimes  prescribed  in  medicine. 

Occurrence,  Preparation,  and  Properties  of  Barium.  —  Barium 
occurs  as  witherite,  the  native  carbonate,  BaCO3,  and  as  barite, 
or  heavy  spar,  the  native  sulphate,  BaSO4.  These  serve  as 
sources  for  the  preparation  of  barium  compounds.  Metallic 


388  OUTLINES   OF  CHEMISTRY 

barium  may  be  obtained  as  a  silver-white  metal  by  electrolysis 
of  the  molten  chloride.  In  general,  its  properties  are  like  those 
of  strontium  and  calcium.  Barium  has  a  specific  gravity  of 
3.78.  It  melts  at  about  850°  and  boils  at  1150°. 

Compounds  of  Barium.  —  These  are  analogous  to  those  of 
calcium  and  strontium.  They  are,  however,  poisonous  in  character. 
Barium  oxide  BaO  is  best  prepared  by  heating  the  nitrate,  for 
the  carbonate  requires  even  a  higher  temperature  to  decompose 
it  than  strontium  carbonate.  On  treating  the  oxide  with  water, 
barium  hydroxide  Ba(OH)2  is  formed,  which  is  more  soluble 
than  calcium  hydroxide.  The  solutions  of  barium  hydroxide 
are  alkaline  and  yield  a  precipitate  of  barium  carbonate  BaCO3 
on  treatment  with  carbon  dioxide.  Barium  hydroxide  is 
frequently  used  in  place  of  lime  water  in  testing  for  carbon 
dioxide. 

On  heating  barium  oxide  in  the  air  or  in  oxygen,  barium 
dioxide  BaO2  is  formed.  This  interesting  substance,  dis- 
covered by  Thenard  and  Gay-Lussac,  is  a  grayish  white  powder 
which  is  insoluble  in  water.  When  barium  dioxide  is  heated 
to  high  temperatures,  it  is  converted  into  barium  oxide  and 
oxygen.  This  affords  a  method  of  preparing  oxygen  from  the 
air,  as  already  mentioned.  Barium  dioxide  is  used  in  prepar- 
ing hydrogen  peroxide.  By  treating  a  solution  of  barium 
hydroxide  with  hydrogen  peroxide  a  precipitate  of  the  compo- 
sition BaO2  •  8  H2O  is  obtained  which  loses  water  at  130°,  and 
yields  barium  dioxide. 

Barium  chloride  BaCl2.  2  H2O,  prepared  by  treating  either  the 
carbonate  or  the  sulphide  with  hydrochloric  acid,  forms  rhombic 
prisms  soluble  in  about  3  parts  of  water.  The  salt  is  not 
deliquescent.  It  is  practically  insoluble  in  alcohol.  This  salt 
is  then  less  soluble  than  the  chloride  of  strontium,  and  the  lat- 
ter is  less  soluble  than  calcium  chloride.  Barium  chloride  has 
a  bitter  taste,  and  is  a  fairly  strong  poison. 

Barium  fluoride  BaF2  is  practically  insoluble  in  water,  but 
readily  soluble  in  acids.  Barium  bromide  BaBr2  and  barium 
iodide  BaI2  are  soluble  in  water,  and  also  in  alcohol. 

Barium  nitrate  Ba(NO3)2  forms  anhydrous  crystals  of  the 
regular  system.  It  is  soluble  in  about  12  parts  of  water,  and 
hence  is  precipitated  by  adding  sodium  nitrate  to  a  concentrated 
solution  of  barium  chloride.  It  may  also  be  obtained  by  the 


THE  ALKALINE  EARTH  METALS  389 

action  of  nitric  acid  on  the  carbonate,  hydroxide,  or  sulphide. 
On  heating,  it  decomposes,  yielding  the  oxide,  oxygen,  and 
oxides  of  nitrogen.  Barium  nitrate  is  used  in  making  green 
Bengal  lights. 

Barium  carbonate  BaCO3  occurs  in  rhombic  crystals  as  with- 
erite,  and  is  prepared  commercially  by  heating  the  natural  sul- 
phate or  barium  sulphide  with  sodium  carbonate.  The  product 
consequently  is  contaminated  with  sodium  carbonate,  which  it 
is  practically  impossible  to  remove  by  washing.  Pure  barium 
carbonate  may  be  obtained  by  adding  ammonium  carbonate  to 
a  solution  of  barium  chloride.  It  requires  white  heat  (1500°  C.) 
to  decompose  barium  carbonate,  but  by  heating  it  with  carbon 
it  may  readily  be  reduced  to  barium  oxide. 

Barium  sulphate  BaSO4  is  the  chief  source  of  barium.  In 
nature  it  is  found  in  compact  masses  or  rhombic  crystals  of 
specific  gravity  4.48.  Barium  sulphate  is  almost  completely 
insoluble  in  water  and  dilute  acids.  Though  alkaline  carbon- 
ates in  aqueous  solution  affect*  barium  sulphate  but  slightly, 
the  latter  may  readily  be  transformed  to  carbonate  by  fusion 
with  sodium  carbonate,  thus  :  — 

BaSO4  -f  Na2CO3  =  Na2SO4  +  BaCO3. 

Barium  sulphate  is  easily  prepared  by  treating  a  solution  of  a 
barium  salt  with  any  soluble  sulphate.  The  precipitate  is  usu- 
ally composed  of  very  small  crystals.  These  grow  larger  on 
standing  and  can  then  be  filtered  off  readily.  In  analytical 
operations,  the  fact  that  barium  sulphate  is  insoluble  and  readily 
prepared  is  much  used  in  testing  for  barium  and  also  for  sul- 
phates. Barium  sulphate  is  often  used  as  a  white  pigment, 
known  as  permanent  white. 

On  heating  barium  sulphate  with  carbon,  barium  sulphide 
BaS  is  produced,  which,  like  the  sulphides  of  calcium  and 
strontium,  phosphoresces  in  the  dark.  This  fact  was  discovered 
in  1603  by  Casciorolo,  a  shoemaker  of  Bologna,  who,  in  an  at- 
tempt to  make  silver,  had  heated  native  barium  sulphate  with 
carbon,  thus  forming  the  sulphide  of  barium.  In  1774  Scheele 
investigated  barium  sulphate  and  showed  that  it  contains  an 
earth  quite  different  from  lime.  Barium  carbonate  was  then 
soon  prepared  artificially,  and  its  probable  occurrence  in  nature 
was  foretold,  when  indeed  in  1783  native  barium  carbonate  was 


390  OUTLINES  OF   CHEMISTRY 

discovered  in  Scotland  by  Withering,  after  whom  the  mineral  is 
called  witherite.  Barium  receives  its  name  from  the  Greek 
word  meaning  heavy.  It  was  reported  that  native  barium  car- 
bonate also  occurred  at  Strontium  in  Scotland.  Investigations 
by  Crawford  in  1790  showed,  however,  that  this  mineral  was 
still  another  carbonate.  Thus  strontium  was  discovered.  It 
derives  its  name  from  the  locality  where  the  native  carbonate, 
strontianite,  was  found. 

Detection  of  the  Alkaline  Earth  Metals.  —  These  may  readily 
be  detected  by  means  of  the  spectroscope,  as  already  stated.  In 
addition  the  following  facts  are  frequently  used  in  detecting 
calcium,  strontium,  and  barium. 

By  adding  a  calcium  sulphate  solution  to  a  solution  of  a  barium 
or  strontium  salt  an  insoluble  sulphate  is  precipitated ;  but  by 
similar  treatment  solutions  of  calcium  salts  remain  clear.  Stron- 
tium sulphate  solution  precipitates  barium  sulphate  from  solu- 
tions of  barium  salts.  Ammonium  oxalate  precipitates  the  oxa- 
lates  of  barium,  strontium,  and  calcium  from  solutions  of  their 
salts.  The  oxalates  of  barium  and  strontium  are  soluble  in  dilute 
acetic  acid,  while  the  oxalate  of  calcium  is  not.  Calcium  nitrate 
is  soluble  in  alcohol,  while  strontium  and  barium  nitrates  are 
not.  Barium  chromate  BaCrO4  is  insoluble  in  water,  while 
calcium  chromate  is  soluble.  The  sulphates  of  calcium  and 
strontium  are  converted  into  carbonates  by  boiling  with  sodium 
carbonate  solution,  while  barium  sulphate  is  not  changed  by 
this  treatment. 

Radium  and  Radio-activity.  —  In  1896  Becquerel  discovered 
that  photographic  plates,  carefully  protected  from  light  by 
means  of  dark  coverings,  are  nevertheless  affected  as  though 
they  had  actually  been  exposed  to  light,  when  kept  near  ura- 
nium or  its  compounds.  Furthermore,  it  was  found  that  a  well- 
insulated,  charged  electroscope  would  soon  become  discharged 
in  the  neighborhood  of  uranium  or  its  compounds,  showing  that 
by  the  latter  the  air  had  been  rendered  a  conductor  of  electric- 
ity. It  was  consequently  concluded  that  these  phenomena  are 
due  to  peculiar  rays  emitted  from  uranium.  These  rays  were 
called  uranium  rays  or  Becquerel  rays,  and  the  substances  from 
which  they  issued  were  called  radio-active.  The  mineral  pitch- 
blende, uraninite,  consisting  essentially  of  a  black  oxide  of 
uranium  U3O8,  together  with  various  impurities  in  minor  quan- 


THE  ALKALINE  EARTH  METALS  391 

titles,  shows  this  radio-activity  in  a  high  degree;  and  M.  and 
Mme.  Curie  found  that  the  residues  of  this  mineral  after  removal 
of  the  uranium  were  still  very  radio-active.  This  led  them  to 
further  investigations,  and  in  1898  they  announced  that  they 
had  discovered  in  these  residues  a  new  element,  radium,  to 
which  the  radio-active  phenomena  are  due.  In  pitchblende 
radium  is  found  together  with  relatively  very  much  larger 
quantities  of  barium,  so  that  a  ton  of  residues  yielded  about 
thirty  pounds  of  barium  chloride,  from  which  by  very  many  re- 
peated fractional  crystallizations  a  few  tenths  of  a  gram  of 
radium  chloride  were  finally  obtained.  This  compound  was 
over  a  million  times  as  radio-active  as  pitchblende. 

Radium  chloride  or  bromide  is  less  soluble  than  the  corre- 
sponding barium  salt,  hence  the  possibility  of  effecting  a  tolerable 
separation  from  the  latter  by  fractional  crystallization.  Radium 
compounds  exhibit  a  spectrum  similar  to  that  of  the  alkaline 
earths.  The  metal  radium  itself  has  not  yet  been  isolated. 

The  analysis  of  the  chloride  has  shown  that  35.46  grams  of 
chlorine  are  combined  with  113.2  grams  of  radium,  hence,  re- 
garding radium  as  bivalent  like  the  alkaline  earth  metals,  to 
which  it  appears  to  be  analogous,  the  atomic  weight  is  226.4. 
The  symbol  of  radium  is  Ra.  On  account  of  the  scarcity  of  the 
element,  but  few  of  its  compounds  have  been  studied.  The 
chloride  RaCl2,  the  bromide  RaBr2,  the  nitrate  Ra(NO3)2,  the 
carbonate  RaCO3,  and  the  hydroxide  Ra(OH)2  have  been  pre- 
pared. All  of  the  compounds  show  radio-activity.  In  the  dark 
they  are  luminous.  They  also  emit  heat  continually.  It  has 
been  estimated  that  a  gram  of  radium  in  its  compounds  gives 
off  heat  at  the  rate  of  approximately  100  calories  per  hour.  As 
the  rays  from  radium  compounds  impinge  upon  a  screen  of  ba- 
rium platinocyanide  or  zinc  sulphide,  they  cause  these  to  phos- 
phoresce; furthermore,  radio-activity  may  be  communicated 
from  radium  compounds  to  substances  placed  near  them.  So 
the  walls  of  containers  of  radio-active  substances  acquire  this 
property,  and  solids  kept  in  contact  with  radium  salts  also 
become  active.  The  rays  emitted  from  radio-active  substances 
are  able  to  kill  germs,  destroy  the  germinating  power  of  seeds, 
and  act  destructively  on  living  tissues,  so  that  experimentation 
with  the  purer  and  consequently  more  powerful  radium 
compounds  must  be  conducted  with  great  care.  Glass  con- 


392  OUTLINES   OF   CHEMISTRY 

tainers  of  radium  salts  slowly  turn  violet  in  color,  while  cloth 
is  gradually  disintegrated  by  the  action  of  the  rays. 

The  rays  emitted  from  compounds  of  radium  are  complex  in 
character.  This  is  shown  by  the  fact  that  a  portion  of  them 
(the  a  rays)  is  readily  absorbed  by  metals,  the  air,  etc.,  and  is 
deflected  with  great  difficulty  by  a  magnetic  field.  Another 
portion  (the  ft  rays)  is  readily  deflected  by  a  magnet,  and  is 
particularly  active  on  photographic  plates;  while  a  third  por- 
tion (the  7  rays)  is  not  deflected  by  a  magnet  and  has  great 
penetrating  power,  passing  even  through  a  thickness  of  30  cm. 
of  iron.  The  study  of  these  rays  has  largely  been  conducted  by 
testing  their  ability  to  affect  photographic  plates  and  to  dis- 
charge an  electroscope. 

Radium  compounds  emit  an  emanation  which  may  be  con- 
densed at  about  —  150°,  by  the  aid  of  liquid  air.  This  emana- 
tion is  therefore  of  the  nature  of  a  gas.  Indeed,  it  may  be  passed 
from  one  tube  to  another  like  air.  When  examined  in  a  vacuum 
tube  by  means  of  the  spectroscope,  the  emanation  from  radium 
compounds  after  a  time  shows  the  spectrum  of  helium.  For  this 
reason  it  has  been  concluded  that  radium  is  slowly  changing  to 
helium.  Furthermore,  since  radium  compounds  gradually  lose 
their  activity,  and  since  uranium  compounds  gradually  again 
acquire  the  ability  of  giving  off  emanations  after  having  been 
deprived  of  the  same,  the  assumption  has  been  made  that 
uranium  is  gradually  changing  into  radium  and  that  the  latter 
is  "  decaying  "  into  emanations,  which  in  turn  are  transformed 
into  helium.  Thus,  the  study  of  radio-active  phenomena  has 
opened  up  the  question  of  the  possibility  of  the  transmutation 
of  the  chemical  elements. 

Thorium  compounds  also  exhibit  radio-activity  and  give  off 
emanations.  A  particularly  active  constituent,  known  as  tho- 
rium Jf,  has  been  separated  by  chemical  means  from  thorium, 
while  a  similar  product,  uranium  X,  has  been  obtained  from 
uranium,  and  another,  actinium  X,  from  actinium.  Actinium 
is  a  radio-active  substance  found  in  uranite  by  Debierne.  It 
acts  like  thorium,  but  more  intensely.  The  elementary  char- 
acter of  actinium  and  of  polonium,  which  has  also  been  found 
in  pitchblende  by  Mme.  Curie,  still  awaits  confirmation. 


CHAPTER  XXIII 

THE   METALS   OF   THE  MAGNESIUM   GROUP 

THE  metals  of  this  group  are  glucinum,  magnesium,  zinc,  cad- 
mium, and  mercury.  Of  these,  magnesium  closely  resembles 
the  alkaline  earth  metals.  Zinc,  cadmium,  and  mercury,  on  the 
other  hand,  have  a  much  higher  specific  gravity,  and  are  much 
less  readily  oxidized;  their  affinity  for  oxygen  diminishes  in  the 
5rder  named.  '  Glucinum,  unlike  the  other  members  of  this 
group,  has  a  high  melting  and  boiling  point.  It  really  is  a 
transition  element  between  the  magnesium  group  and  the  metals 
of  the  earths. 

Glucinum,  or  Beryllium,  Gl,  or  Be,  At.  Wt.  9.1,  is  found  in 
nature  as  a  constituent  of  the  rather  rare  minerals  beryl 
Gl3Al2Si6O18,  .phenacite  Gl2SiO4,  and  chrysoberyl  G1(A1O2)2. 
When  colored  green  by  traces  of  chromium  silicate,  beryl  is 
called  emerald,  which  is  used  as  a  gem.  Glucinum  is  a  white, 
malleable  metal  having  a  specific  gravity  1.8.  It  does  not 
liberate  hydrogen  from  water  even  on  boiling.  It  dissolves  in 
hydrochloric  or  sulphuric  acid,  while  nitric  acid  attacks  it  but 
slightly.  In  caustic  potash,  glucinum  dissolves  readily,  liberat- 
ing hydrogen.  The  metal  is  obtained  by  electrolysis  of  molten 
chrysoberyl,  or  by  heating  the  oxide  with  magnesium  powder. 
In  its  compounds,  glucinum  is  always  bivalent.  The  compounds 
have  a  sweetish  taste,  whence  the  name  glucinum  (sweet);  the 
element  is  also  called  beryllium  from  beryl,  in  which  it  occurs. 

Glucinum  oxide  G1O  was  discovered  in  1797  by  Vauquelin. 
It  is  basic,  but  also  possesses  very  weak  acidic  properties,  form- 
ing soluble  salts  with  caustic  alkalies.  From  solutions  of 
glucinum  salts,  ammonium  hydroxide  precipitates  glucinum 
hydroxide  G1(OH)2. 

Among  the  well-known  salts  are  glucinum  sulphate  G1SO4, 
glucinum  chloride  G1C12,  and  glucinum  carbonate  G1CO3.  Basic 
salts,  notably  basic  nitrates  and  sulphates,  are  also  known. 

Occurrence,  Preparation,  and  Properties  of  Magnesium. — Mag- 
nesium is  very  widely  distributed.  It  occurs  in  large  quanti- 

393 


394  OUTLINES  OF   CHEMISTRY 

ties  in  dolomite  MgCO3-CaCO3,  which  often  forms  mountains. 
It  is  also  found  in  carnallite  MgCl2-KCl-6H2O,  in  kieserite 
MgSO4-H2O,  in  schonite  MgSO4  •  K2SO4  -  6  H2O,  in  kainite 
K2SO4-MgSO4-  MgCl2-  6  H2O,  in  magnesite  MgCO3,  in  soapstone 
or  talc  Mg3H2Si4O12,  in  meerschaum  Mg2Si3O8-4  H2O,  in  serpen- 
tine Mg3Si2O7  •  2  H2O,  in  hornblende  Mg2CaFeSi4O12,  in  asbes- 
tus  Mg3Si2O7  •  2  H2O,  and  in  many  other  complex  silicates.  Sea 
water  contains  magnesium  chloride  and  sulphate.  These  salts 
also  occur  in  many  spring  waters,  which  are  called  bitter  waters. 
Magnesium  salts  are  found  in  soils,  being  decomposition  prod- 
ucts of  rocks.  Plants  and  animals  contain  magnesium,  which 
appears  in  their  ash  as  carbonate  and  phosphate.  Guano  con- 
tains magnesium  ammonium  phosphate. 

Metallic  magnesium  is  prepared  by  electrolyzing  molten  carnal- 
lite  in  an  iron  pot,  which  also  serves  as  cathode,  a  carbon  anode 
being  employed.  The  metal  is  silver-white,  and  when  hot  it 
may  be  drawn  into  ribbon  or  wire.  At  ordinary  temperatures 
magnesium  is  but  slightly  malleable.  Exposed  to  air,  it  grad- 
ually becomes  covered  with  a  thin,  white  layer  of  oxide.  It 
melts  at  633°  and  boils  at  1100.°  Its  specific  gravity  is  1.75. 
It  may  be  sublimed  in  a  vacuum,  yielding  hexagonal,  prismatic 
crystals.  In  an  atmosphere  of  hydrogen  the  metal  may  be  dis- 
tilled. It  acts  but  very  slightly  on  pure  water  even  at  100° ; 
but  when  heated  in  a  current  of  steam,  it  takes  fire.  Dilute 
acids  act  vigorously  on  magnesium.  In  the  air,  the  metal  burns 
with  a  brilliant  white  light,  yielding  the  oxide  MgO,  together 
with  some  nitride  Mg3N2.  Magnesium  light  has  a  powerful 
effect  upon  photographic  plates  and  is  consequently  used  in  flash 
lights.  Flash  light  powder  consists  of  about  5  parts  of  powdered 
magnesium  to  9  parts  of  potassium  chlorate.  Magnesium  is 
also  used  in  fireworks,  for  signal  lights,  and  as  a  reducing  agent 
in  chemical  operations. 

The  metal  wras  first  obtained  in  pure  form  by  Liebig  and 
Bussy  in  1830.  Bunsen  prepared  it  by  electrolysis  twenty-two 
years  later.  The  atomic  weight  of  magnesium  is  24.32,  and  its 
valence  is  two. 

Magnesium  Oxide  MgO  is  a  white  powder,  usually  prepared 
by  heating  the  carbonate.  It  is  also  called  magnesia,  magnesia 
usta,  or  calcined  magnesia.  In  contact  with  water  it  does  not 
dissolve,  but  it  forms  magnesium  hydroxide  Mg(OH)2,  which  is 


THE  METALS  OF  THE  MAGNESIUM  GROUP 


395 


very  sparingly  soluble.  Magnesium  oxide  is  used  in  medicine. 
It  is  even  more  difficult  to  melt  magnesia  than  lime,  conse- 
quently magnesia  is  used  in  making  fire  brick  and  in  the  con- 
struction of  the  electric  furnace.  Magnesium  oxide  does  not 
have  acidic  properties. 

Magnesium  Carbonate  MgCO3  occurs  in  nature  as  magnesite  in 
hexagonal  crystals  that  are  isomorphous  with  calcite.  On  adding 
a  solution  of  an  alkaline  carbonate  to  a  solution  of  a  magnesium 
salt,  basic  magnesium  carbonate  is  precipitated,  the  composition 
of  which  varies  according  to  the  temperature  and  concentration 
of  the  solutions  used.  The  ordinary  magnesium  carbonate  of 
commerce,  called  magnesia  alba,  is  prepared  by  precipitation. 
Its  composition  is  approximately  (MgCO3)3-Mg(OH)2-  3  H2O. 
Though  insoluble  in  water,  magnesium  carbonate,  like  calcium 
carbonate,  dissolves  in  water  charged  with  carbon  dioxide. 
From  such  solutions  crystals  of  the  composition  MgCO3-3  H2O 
(and  at  low  temperatures  MgCO3-5H2O)  have  been  obtained. 
Magnesium  carbonate  is  used  in  medicine,  also  as  a  cosmetic. 

Magnesium  Chloride  MgCl2  •  6  H2O  forms  highly  hygroscopic, 
monoclinic  crystals.  On  heating  these  they  completely  decom- 
pose, thus :  — 

MgCl2  -  6  H20  =  MgO  +  2  HC1  +  5  H2O. 

The  anhydrous  salt  MgCl2  may  be  obtained  by  heating  magne- 
sium ammonium  chloride  MgCl2-NH4Cl-6  H2O.  Figure  136 


40  60  80  100 

Temperature  in  degrees  C. 
FIG.  136. 


120         140         160         180 


396  OUTLINES  OF  CHEMISTRY 

presents  the  solubility  curve  of  magnesium  chloride.  It  will  bo 
seen  from  the  figure  that  five  different  hydrates  are  known. 

Magnesium  Sulphate  MgSO4  •  7  H2O,  also  known  as  Epsom 
salt  and  bitter  salt,  forms  large  rhombic  prisms  that  dissolve 
in  about  four  parts  of  water.  The  solution  has  a  disagreeable 
bitter  taste.  The  salt  is  found  in  the  waters  of  Epsom  springs 
and  many  other  mineral  springs.  It  is  used  as  a  purgative,  and 
is  also  employed  in  "  loading  "  cotton  goods  and  in  making  the 
sulphates  of  potassium  and  sodium.  Occasionally  it  serves  as  a 
fertilizer. 

Magnesium  Phosphates  are  in  general  similar  to  those  of 
calcium.  Magnesium  ammonium  phosphate  MgNH4PO4  is  im- 
portant in  analytical  chemistry.  The  salt  forms  rhombic 
crystals  that  are  insoluble  in  ammonia  water.  On  ignition  they 
yield  magnesium  pyrophosphate  Mg2P2O7. 

Magnesium  Ammonium  Arsenate  MgNH4AsO4  is  analogous 
to  magnesium  ammonium  phosphate.  On  ignition  it  yields 
magnesium  pyroarsenate  Mg2As2O7. 

Tests  for  Magnesium.  —  In  testing  for  magnesium  compounds, 
the  following  facts  are  of  importance  :  — 

Basic  carbonate  of  magnesium  is  precipitated  by  carbon- 
ates of  the  alkalies.  Hydroxides  of  the  alkalies  precipitate 
magnesium  hydroxide ;  but  ammonium  hydroxide  does  not 
precipitate  magnesium  hydroxide  in  presence  of  ammonium 
chloride,  for .  the  solutions  of  the  latter  dissolve  magnesium 
hydroxide.  In  presence  of  ammonium  chloride,  clear  am- 
moniacal  solutions  of  magnesium  salts  are  precipitated  by 
sodium  phosphate,  magnesium  ammonium  phosphate  being 
formed. 

Occurrence,  Preparation,  and  Properties  of  Zinc.  —  Zinc  occurs 
in  nature  mainly  as  the  carbonate,  ZnCO3,  in  the  mineral  called 
smithsonite,  calamine  or  zinc  spar,  and  as  the  sulphide,  ZnS,  in 
zinc  blende  or  black-jack.  Other  ores  of  zinc  are  franklinite 
Zn(FeO2)2,  gahnite  Zn(AlO2)2,  also  called  zinc  spinel,  and  red 
zinc  ore  ZnO.  Most  of  the  ores  of  zinc  also  contain  some 
cadmium. 

In  making  metallic  zinc,  the-  ores  are  first  roasted.  Thus 
carbonates ,  and  sulphides  are  finally  all  transformed  to  zinc 
oxide.  The  latter  is  then  reduced  by  heating  it  with  carbon. 
The  reactions  involved  are  :  — 


THE   METALS   OF  THE   MAGNESIUM  GROUP  397 

ZnCO3  =  ZnO  +  CO2. 
ZnS  +  3  O  =  ZnO  +  SO2. 
ZnO  +  C  =  Zn  +  CO. 

The  operation  of  reducing  the  zinc  oxide  is  conducted  in 
earthenware  retorts.  At  the  bright  red  heat  developed,  1200° 
to  1300°,  zinc  is  converted  into  vapor,  which  is  condensed  in 
iron  receivers  connected  with  the  retorts.  Zinc  dust,  consist- 
ing of  finely  divided  zinc  plus  5  to  10  per  cent  of  zinc  oxide,  is 
at  first  obtained  on  the  sides  of  the  condenser  ;  later,  the  metal 
condenses  to  a  liquid  and  is  run  into  molds.  This  crude  zinc, 
called  spelter,  is  contaminated  with  carbon,  iron,  arsenic,  lead, 
and  cadmium.  It  is  purified  by  redistillation.  Chemically 
pure  zinc  is  made  by  heating  pure,  precipitated  zinc  carbonate 
with  pure  carbon,  or  still  better  by  the  electrolysis  of  pure  zinc 
salts. 

Zinc  crystallizes  in  the  hexagonal  system.  It  is  bluish  white, 
and  has  a  bright  metallic  luster.  Cast  zinc  has  a  specific 
gravity  of  6.9  ;  but  hammered  zinc,  or  zinc  wire,  is  denser, 
having  a  specific  gravity  of  7.2.  Zinc  is  rather  brittle,  but  at 
120°  it  becomes  malleable  and  ductile.  When  heated  to  200°, 
it  becomes  so  brittle  that  it  may  be  pulverized.  Zinc  melts  at 
420°  and  boils  at  918°.  Its  atomic  weight  is  65.37  and  its  molec- 
ular weight  is  the  same,  for  the  vapor  is  33.93  times  as  heavy 
as  hydrogen.  Zinc  is  always  bivalent.  On  exposure  to  moist 
air,  the  luster  of  zinc  is  dimmed  by  the  formation  of  a  thin 
coating  of  white  basic  carbonate.  At  high  temperatures,  zinc 
burns  in  the  air  with  a  brilliant,  bluish  white  flame.  On  water 
zinc  does  not  act,  but  when  steam  is  passed  over  heated  zinc, 
the  oxide  and  hydrogen  are  formed.  Nearly  all  dilute  acids 
act  on  zinc ;  the  purer  the  zinc,  the  less  rapid  is  the  action.  In 
general,  hydrogen  is  liberated  when  zinc  acts  on  acids,  though 
in  some  cases,  like  that  of  nitric  acid  for  instance,  the  hydrogen 
is  not  set  free,  for  it  at  once  reduces  some  of  the  acid  present. 
On  heating  zinc  with  concentrated  sulphuric  acid,  sulphur 
dioxide  and  zinc  sulphate  are  formed  :  — 

Zn  +  2  H2SO4  =  ZnSO4  +  SO2  +  2  H2O. 

In  hot  caustic  alkalies,  zinc  dissolves,  as  already  stated,  forming 
a  zincate  and  hydrogen  :  — 

Zn  +  2  NaOH  =  NaaZnOa  +H2. 


398  OUTLINES   OF  CHEMISTRY 

When  introduced  into  solutions  of  many  of  the  salts  of  either 
lead,  tin,  copper,  mercury,  silver,  platinum,  or  gold,  zinc  precipi- 
tates these  metals,  usually  as  finely  divided  powders,  thus  :  — 

CuSO4  +  Zn  =  ZnSO4  +  Cu. 

Zinc  is  used  in  sheet  form  for  many  purposes,  such  as  making 
roofs,  gutters,  and  architectural  ornaments.  Galvanized  iron, 
so-called,  consists  of  sheet  iron  which  has  been  coated  with  zinc 
by  dipping  the  thoroughly  clean  sheets  of  iron  into  highly 
heated  molten  zinc.  The  zinc  coating  prevents  the  iron  from 
rusting.  Much  zinc  is  also  used  in  making  electrical  batteries 
(which  see),  and  in  preparing  many  alloys,  especially  brass. 

Brass,  which  is  an  alloy  of  zinc  and  copper,  was  known  long 
before  zinc,  for  it  was  obtained  by  melting  native  copper  ores 
that  contained  zinc.  On  the  European  continent  zinc  produc- 
tion on  a  commercial  scale  began  about  one  hundred  years  ago  ; 
though  in  England  zinc  was  manufactured  some  fifty  years 
earlier.  In  1910  the  United  States  produced  252,497  tons  of  zinc, 
which  is  approximately  one  third  of  the  world's  annual  output. 

Zinc  Oxide  ZnO,  also  called  flores  zinci  or  lana  philosophica, 
is  a  white,  bulky  powder  obtained  by  burning  zinc  in  the  air  or 
by  heating  the  basic  carbonate  of  zinc.  When  hot,  it  is  yellow, 
but  on  cooling,  it  turns  white.  It  is  much  used  in  white  paints 
under  the  name  zinc  white.  The  oxide  is  also  used  in  pharmacy 
for  making  ointments.  Zinc  hydroxide  is  obtained  as  an  amor- 
phous precipitate  by  adding  caustic  alkalies  to  solutions  of  zinc 
salts.  The  hydroxide  is  soluble  in  an  excess  of  the  precipitant, 
forming  a  zincate.  On  heating  the  hydroxide,  it  loses  water, 
forming  the  oxide.  Native  oxide  of  zinc  is  frequently  colored 
red  because  of  the  presence  of  oxides  of  manganese. 

Zinc  Carbonate  ZnCO3  occurs  in  nature,  as  already  stated. 
It  forms  rhombohedral  crystals  that  are  isomorphous  with  cal- 
cite.  On  treating  solutions  of  zinc  salts  with  alkali  carbonates, 
basic  zinc  carbonates  are  precipitated,  whose  composition  varies 
according  to  the  temperature  and  concentration  of  the  solu- 
tions. These  carbonates  are  approximately  ZnCO3  •  2  Zn(OH)2 
or  (ZnCO3)2  •  3  Zn(OH)2.  They  are  soluble  in  an  excess  of  am- 
monium carbonate,  but  not  in  sodium  or  potassium  carbonate. 

Zinc  Chloride  ZnCl2  is  made  by  burning  zinc  in  chlorine,  or 
by  the  action  of  hydrochloric  acid  on  the  carbonate,  oxide,  hy- 


THE  METALS  OF  THE  MAGNESIUM  GROUP  39§ 

dioxide,  or  the  metal  itself.  It  is  a  white,  deliquescent  mass, 
having  caustic  properties.  It  is  soluble  in  alcohol  as  well  as 
in  water.  On  attempting  to  dehydrate  the  salt  by  heating  it, 
hydrochloric  acid  is  given  off,  as  in  the  case  of  magnesium  chlo- 
ride. On  mixing  concentrated  zinc  chloride  solutions  with  zinc 
oxide,  oxychlorides  like  Zn(OH)Cl  are  formed,  the  mass  hard- 
ening in  a  manner  not  unlike  the  setting  of  plaster  of  Paris. 
Such  zinc  chloride  and  zinc  oxide  mixtures  are  used  as  cements 
in  filling  teeth.  Zinc  chloride  is  further  employed  in  preserv- 
ing railroad  ties,  which,  after  being  soaked  in  its  solutions,  do 
not  rot  readily.  It  is  also  used  to  cleanse  the  surface  of  metals 
in  the  process  of  soldering.  In  medicine  zinc  chloride  is  used 
as  a  caustic  agent  and  as  a  disinfectant.  In  the  laboratory  it 
is  used  in  connection  with  certain  syntheses  of  organic  com- 
pounds. Zinc  bromide  ZnBr2  is  analogous  to  the  chloride.  The 
iodide  ZnI2  is  also  very  soluble ;  it  readily  splits  off  iodine. 
The  fluoride  ZnF2  is  but  sparingly  soluble. 

Zinc  Sulphate  ZnSO4,  white  vitriol,  is  formed  by  the  oxida- 
tion of  zinc  sulphide  or  by  the  action  of  sulphuric  acid  on  the 
metal,  oxide,  or  carbonate.  It  crystallizes  in  the  rhombic  sys- 
tem in  colorless  prisms  of  the  composition  ZnSO4  •  7  H2O,  which 
effloresce  on  exposure  to  the  air.  The  crystals  are  ismorphous 
with  MgSO4  •  7  H2O.  They  are  readily  soluble  in  water. 
With  sulphates  of  the  alkalies,  zinc  sulphate  forms  double 
salts,  like  K2SO4  •  ZnSO4  •  6  H2O,  which  is  analogous  to  scho- 
nite,  K2SO4  •  MgSO4  •  6  H2O.  Like  other  zinc  compounds,  zinc 
sulphate  is  moderately  poisonous,  upon  which  its  use  as  an 
antiseptic  depends. 

Zinc  Sulphide  ZnS  is  white  when  pure.  Native  sulphide  of 
zinc,  black-jack,  is  colored  dark  brown  by  feme  oxide  and 
other  impurities.  It  crystallizes  in  the  regular  system.  On 
adding  sulphides  of  the  alkalies  or  of  ammonium  to  a  solution 
of  a  zinc  salt,  a  white  precipitate  of  zinc  sulphide  is  formed. 
This  is  soluble  in  dilute  mineral  acids  but  not  in  acetic  acid. 

Analytical  Tests  for  Zinc  Salts.  —  The  reactions  that  are  used 
in  testing  for  zinc  salts  are  :  — 

Ammonium  sulphide  precipitates  white  zinc  sulphide  from 
solutions  of  zinc  salts.  The  sulphide  is  insoluble  in  acetic  acid 
but  soluble  in  mineral  acids. 

Alkaline  hydroxides  precipitate  zinc  hydroxide  from  solutions 


400  OUTLINES   OF  CHEMISTRY 

of  zinc  salts.  The  precipitate  dissolves  in  an  excess  of  the 
reagent. 

Alkaline  carbonates  precipitate  basic  zinc  carbonate,  which 
is  soluble  in  ammonium  carbonate  but  not  in  sodium  or  potas- 
sium carbonate. 

Zinc  oxide  is  white  when  'cold  and  yellow  when  hot.  When 
moistened  with  cobalt  nitrate  solution  and  then  strongly  heated 
on  charcoal,  zinc  oxide  yields  a  green  mass  called  Rinmann's 
green. 

Occurrence,  Preparation,  and  Properties  of  Cadmium. — Cadmium 
is  a  rather  rare  element.  It  frequently  occurs  in  small  quanti- 
ties in  ores  of  zinc.  It  is  also  found,  though  rarely,  as  the 
native  sulphide,  CdS,  known  as  greenockite,  which  is  hexagonal 
and  isomorphous  with  a  rare,  native  zinc  sulphide,  wiirzite. 
Cadmium  is  commonly  obtained  in  connection  with  the  reduc- 
tion of  zinc  from  its  ores.  The  metal  boils  at  775°  and  hence 
passes  over  into  the  condensers  before  the  zinc,  which  boils  at 
918°.  Cadmium  is  silver-white,  malleable  and  ductile  even  at 
ordinary  temperatures.  It  resembles  tin  in  appearance,  melts 
at  320°,  and  has  a  specific  gravity  of  8.6.  Its  atomic  weight  is 
112.4,  and  its  molecular  weight  is  the  same,  for  its  vapor  is 
55.65  times  as  heavy  as  hydrogen.  In  the  air,  cadmium  is  quite 
stable,  though  its  luster  becomes  somewhat  tarnished,  due  to  the 
formation  of  a  thin  layer  of  oxide.  When  strongly  heated  in 
the  air  or  in  oxygen,  it  burns,  forming  a  brown  oxide.  Cadmium 
is  always  bivalent.  Dilute  acids  act  on  cadmium,  liberating 
hydrogen.  The  element  was  discovered  in  1817  by  Stromeyer, 
who  investigated  a  yellowish  zinc  oxide  which  contained  no 
iron.  Almost  simultaneously,  cadmium  was  discovered  by 
Hermann.  The  metal  is  used  in  preparing  standard  cells  for 
measuring  electromotive  forces.  It  is  also  used  in  certain 
alloys  of  low  melting  point,  like  Wood's  metal,  which  has 
already  been  described. 

Cadmium  Compounds.  —  The  oxide,  CdO,  is  a  brown  powder 
which  is  readily  reduced  at  higher  temperatures  by  means  of 
carbon  or  hydrogen.  The  hydroxide,  Cd(OH)2,  is  white.  It  is 
formed  by  leaving  the  oxide  in  contact  with  water  or  by  pre- 
cipitation from  solutions.  It  is  soluble  in  ammonium  hydroxide. 
The  chloride,  CdCl2,  may  be  obtained  from  solutions  in  the  form 
of  crystals  of  the  composition  CdCl2  •  2  H2O,  which  effloresce. 


THE  METALS  OF  THE  MAGNESIUM  GROUP       401 

The  anhydrous  salt  melts  at  540°,  and  may  be  distilled  at 
about  900°.  The  bromide,  CdBr2  •  4  H2O,  is  soluble  in  alcohol 
as  well  as  in  water.  The  iodide,  CdI2,  is  obtained  by  the  direct 
union  of  iodine  and  cadmium  in  presence  of  water.  This  salt 
is  soluble  in  alcohol,  and  is  used  in  photography.  The  nitrate, 
Cd(NO3)2  •  4  H2O,  is  a  deliquescent  salt  obtained  by  the  action 
of  nitric  acid  upon  either  the  metal,  the  hydroxide,  or  the  car- 
bonate. The  sulphate,  3  CdSO4  •  8  H2O,  forms  rnonoclinic  crys- 
tals that  readily  dissolve  in  water.  The  salt  effloresces  on 
exposure  to  the  air.  It  will  be  observed  that  the  sulphate  is 
not  analogous  to  zinc  sulphate  ZnSO4  •  7  H2O  and  magnesium 
sulphate  MgSO4  •  7  H2O.  Cadmium  sulphate  is  used  in  treat- 
ing diseases  of  the  eye.  The  sulphide,  CdS,  is  obtained  as  a 
bright  yellow  precipitate  by  conducting  hydrogen  sulphide  into 
a  solution  of  a  cadmium  salt.  Dilute  acids  do  not  affect  it. 
It  is  used  as  a  pigment  in  paints. 

The  behavior  of  the  hydroxide,  carbonate,  and  sulphide,  as 
above  described,  is  used  in  testing  for  cadmium  salts. 

Occurrence,  Preparation,  and  Properties  of  Mercury.  —  Mercury, 
or  hydrargyrum,  meaning  silver  water,  is  the  only  metal  that 
is  liquid  at  ordinary  temperatures.  It  ,was  known  at  least 
three  hundred  years  before  the  Christian  era.  In  nature  it  is 
sometimes  found  uncombined  in  small  drops  in  the  interstices 
of  rocks.  Its  principal  ore  is  cinnabar,  the  sulphide,  HgS, 
which  forms  dark  red,  hexagonal,  prismatic  crystals.  Mercury 
ores  occur  at  Almaden  in  .Spain,  Idria  in  Austria,  New  Alma- 
den  in  California,  in  Prussia,  Peru,  Japan,  and  China.  The 
metal  is  obtained  by  simply  roasting  the  sulphide  :  — 


The  mercury  vapors  are  condensed  and  collected.  Sometimes 
the  ore  is  heated  with  lime  in  iron  retorts,  when  calcium  sul- 
phide remains  behind  and  mercury  passes  over  in  form  of  vapor 
and  is  condensed.  Mercury  is  purified  by  redistillation  after 
treatment  with  dilute  nitric  acid,  ferric  chloride  solution,*  or 
dilute  sulphuric  acid  plus  potassium  bichromate.  Mechanical 
impurities  are  removed  from  mercury  by  filtration  through 
cloth  or  chamois  skin.  Perfectly  pure  mercury  is  made  by  lib- 
erating the  metal  from  pure  mercury  salts,  and  finally  distilling 
the  product,*  preferably  in  vacuo. 

2D 


402  OUTLINES   OF  CHEMISTRY 

Mercury  is  silver-white,  and  has  a  brilliant  metallic  luster. 
It  melts  at  —39.4°  and  boils  at  357°.  Its  specific  gravity  is 
13.59.  Its  atomic  weight  is  200.6,  which  is  also  its  molecular 
weight,  for  mercury  vapor  is  about  100  times  as  heavy  as 
hydrogen.  Mercury  vapor  has  a  very  characteristic  spectrum, 
exhibiting  a  bright  yellow  and  a  green  line,  besides  a  red, 
a  blue,  and  three  violet  lines.  The  vapor  conducts  electricity 
and  emits  a  bright,  pale,  greenish  light  rich  in  rays  that  affect 
photographic  plates.  The  mercury  lamp  depends  on  the  prin- 
ciple that  mercury  vapors  emit  light  when  they  are  conducting 
high  tension  electricity. 

The  vapors  of  mercury  are  very  poisonous,  which  is  also  true 
of  compounds  of  mercury.  Mercury  crystallizes  in  the  regular 
system.  Solid  mercury  is  malleable.  It  may  be  cut  with  tools, 
and  beaten  into  sheets  with  a  hammer.  In  the  air  mercury 
remains  practically  unchanged.  By  sulphuric  or  hydrochloric 
acids  it  is  attacked  but  slightly.  Nitric  acid  or  hot,  concen- 
trated sulphuric  acid  readily  dissolves  it.  With  sulphur  or  the 
halogens  it  combines  at  slightly  elevated  temperatures. 

Mercury  is  used  in  making  thermometers,  barometers,  mir- 
rors, various  amalgams  in  dentistry,  and  many  compounds  that 
find  application  in  medicine.  Large  amounts  of  mercury  are 
also  employed  in  extracting  gold  and  silver 
from  their  ores.  In  1910  the  United  States 
produced  773  tons  of  mercury,  which  is  about 
one  fourth  of  the  total  amount  produced  that 
year  in  the  world.  Mercury  is  shipped  in 
iron  flasks  containing  about  75  pounds  each 
(Fig.  137). 

Amalgams.  —  The  alloys  or  combinations  of 
mercury  with  other  metals  are  called  amal- 
gams. These  are  usually  made  by  simply 
bringing  the  metals  in  contact  with  mercury, 
though  they  may  also  be  obtained  by  electro- 
lyzing  a  salt,  using  a  mercury  cathode ;  or 
frequently  by  introducing  a  metal  into  a  solu- 
tion of  a  mercury  salt,  like  the  nitrate.  In 
IG'  '  general,  amalgams  are  of  the  nature  of  solutions 

of  the  metals  in  mercury.     They  are  liquid  when  the  mercury 
preponderates,  and  solid  when  relatively  less  mercury  is  used. 


THE   METALS   OF  THE   MAGNESIUM   GROUP  403 

Sodium  combines  with  mercury  with  evolution  of  light  and 
heat.  Sodium  amalgam  containing  less  than  1  per  cent 
sodium  is  liquid ;  a  1  per  cent  sodium  amalgam  is  viscous, 
while  amalgams  containing  2  per  cent  or  more  of  sodium 
are  solid.  Crystalline  sodium  amalgams  corresponding  to 
the  formulae  Na3Hg  and  NaHg6  have  been  isolated.  Sodium 
amalgam  is  used  as  a  reducing  agent,  as  already  mentioned 
under  sodium.  With  the  exception  of  iron  and  platinum, 
practically  all  the  metals  form  amalgams  with  mercury.  Even 
platinum  may  be  amalgamated  by  electrolyzing  a  mercury 
salt  with  a  platinum  cathode.  However,  the  readiness  with 
which  the  various  metals  unite  with  mercury  varies  very 
greatly.  So  gold  and  silver  dissolve  with  ease  in  mercury, 
which  fact  is  used  in  extracting  these  metals  from  pulver- 
ized rocks  with  which  they  are  mixed.  Copper,  cadmium, 
and  tin  also  readily  unite  with  mercury.  The  amalgam  used 
in  filling  teeth  usually  consists  of  mercury  mixed  with  an 
alloy  containing  essentially  silver  and  tin,  together  with  smaller 
amounts  of  other  metals,  among  which  are  copper,  cadmium, 
and  gold.  The  amalgam  is  made  into  a  stiff  paste  and  intro- 
duced into  the  cavity  in  the  tooth,  where  it  soon  hardens  or 
sets  without  material  change  of  volume.  Tin  amalgam  is  used 
in  making  mirrors.  Amalgamated  zinc  is  hardly  acted  upon  by 
dilute  acids;  on  the  other  hand,  magnesium  amalgam,  which 
forms  only  on  heating  mercury  and  magnesium  together,  is  a 
black  incoherent  mass  which  violently  decomposes  water  with 
liberation  of  hydrogen. 

When  sodium  amalgam  is  treated  with  a  concentrated  solu- 
tion of  ammonium  chloride  or  other  ammonium  salt,  a  bulky 
mass  called  ammonium  amalgam  is  obtained.  It  is  supposed  to 
contain  ammonium  NH4,  dissolved  in  mercury.  The  large  bulk 
is  produced  by  hydrogen  and  ammonia  gases  that  are  set  free. 
This  so-called  ammonium  amalgam  is  also  formed  by  electro- 
lyzing a  solution  of  an  ammonium  salt,  using  mercury  as  a 
cathode.  By  some  it  is  considered  practically  certain  that  the 
material  contains  ^H^,  dissolved  in  mercury  ;  by  others  this  has 
been  denied.  At  any  rate,  if  ammonium  amalgam  exists  at  all, 
it  is  very  unstable. 

Compounds  of  Mercury.  —  Mercury  forms  two  distinct  series 
of  compounds :  the  mercurous  compounds,  in  which  the  metal 


404  OUTLINES   OF  CHEMISTRY 

is  univalent  ;  and  the  mercuric  compounds,  in  which  it  is 
bivalent. 

Oxides  of  Mercury.  —  Mercurous  oxide  Hg2O  is  a  dark  brown, 
slightly  greenish,  unstable  powder  obtained  as  a  precipitate  by 
adding  sodium  or  potassium  hydroxide  solution  to  a  mercurous 
salt.  It  decomposes  into  mercuric  oxide  and  mercury  on  ex- 
posure to  light  or  when  slightly  heated.  Mercuric  oxide  HgO 
may  be  obtained  by  prolonged  heating  of  mercury  in  the  air  at 
about  360°.  It  is  made  on  a  large  scale  by  heating  a  mixture 
of  mercuric  nitrate  and  mercury.  It  is  a  bright  red,  crystalline 
powder  also  known  as  red  precipitate.  On  heating  mercuric 
oxide,  it  turns  dark  and  gives  off  oxygen.  Above  500°  the  de- 
composition into  mercury  and  oxygen  progresses  rapidly.  Red 
oxide  of  mercury  is  used  in  medicine  in  red  precipitate  oint- 
ment. The  yellow  variety  is  also  used  for  similar  purposes  ; 
being  finely  divided,  it  is  more  active  than  the  red.  By  adding 
caustic  soda  to  a  solution  of  a  mercuric  salt,  mercuric  oxide  is 
obtained  in  form  of  a  finely  divided,  amorphous,  yellow  powder. 
On  heating,  it  behaves  like  the  red  oxide.  Hydroxides  of 
mercury  are  not  known,  probably  because  they  are  so  unstable, 
decomposing  at  once  into  oxide  and  water. 

Halides  of  Mercury.  —  Mercurous  chloride  HgCl,  also  called 
calomel,  is  obtained  as  a  slightly  yellowish,  crystalline  mass  by 
heating  mercury  with  mercuric  chloride  :  — 


It  may  also  be  formed  by  passing  sulphur  dioxide  into  a  hot 
solution  of  mercuric  chloride  :  — 

2HgCl2  +  2H2O  +  SO2  =  2  HC1  +  H2SO4  -f  2HgCl. 

Other  reducing   agents  may  be  employed  instead  of   sulphur 
dioxide;  thus,  with  stannous  chloride  the  reaction  is, 
2  HgCl2  +  SnCl2  =  SnCl4  +  2  HgCl. 

By  adding  a  soluble  chloride  to  a  solution  of  a  mercurous  salt, 
mercurous  chloride  is  precipitated,  thus  :  — 

HgNO3  +  NaCl  =  NaNO3  +  HgCl. 

Mercurous  chloride  is  not  soluble  in  water.  It  may  be  sub- 
limed, and  thus  obtained  in  crystalline  form.  The  precipitated 
variety  is  an  amorphous  powder.  This  salt  is  much  used  in  medi- 
cine as  a  purgative  and  as  a  stimulant  for  the  secretory  organs. 


THE   METALS   OF   THE   MAGNESIUM   GROUP  405 

It  should  be  kept  in  the  dark,  for  on  exposure  to  light  it  gradu- 
ally decomposes  somewhat,  yielding  mercury  and  mercuric  chloride, 
which  is  a  strong  poison,  thus :  — 

2HgCl  =  Hg  +  HgCl2. 

Mercurous  bromide  HgBr  is  obtained  as  a  white,  insoluble 
powder  by  adding  a  soluble  bromide  to  a  mercurous  salt.  It 
may  be  obtained  in  crystalline  form  by  sublimation,  or  by 
treating  mercury  with  bromine  water. 

Mercurous  iodide  Hgl  is  formed  by  triturating  mercury  and 
iodine  together  in  presence  of  a  little  alcohol,  or  by  precipitat- 
ing a  solution  of  a  mercurous  salt  with  sodium  or  potassium 
iodide.  It  is  a  dark,  yellowish  green  powder  which  readily 
decomposes  into  mercury  and  mercuric  iodide,  especially  in  the 
light,  thus  :  — 

2  Hgl  =  Hg  +  HgI2. 

The  salt  is  used  In  medicine. 

Mercuric  chloride  HgCl2,  also  called  corrosive  sublimate,  or 
sublimate,  is  made  on  a  large  scale  by  heating  mercuric  sul- 
phate with  common  salt,  thus  :  — 

HgS04  +  2  NaCl  =  Na2SO4  +  HgCl2. 

In  this  process  the  mercuric  chloride  vapors  formed  are  con- 
densed, thus  yielding  beautiful,  rhombic,  prismatic  crystals  that 
readily  dissolve  in  water  (1  in  15)  and  also  in  alcohol  and  in 
ether.  The  salt  melts  at  265°  and  boils  at  307°.  Mercuric 
chloride  may  also  be  made  by  the  action  of  hydrochloric  acid 
on  mercuric  oxide,  or  by  dissolving  mercury  in  aqua  regia. 
Mercuric  chloride  is  reduced  to  mercurous  chloride  by  reduc- 
ing agents  as  already  stated  above.  An  excess  of  stannous 
chloride  reduces  mercuric  chloride  to  metallic  mercury :  — 

HgCl2  +  SnCl2=  SnCl4  +  Hg. 

Mercuric  chloride  is  a  powerful  poison.  It  is  used  as  a  disin- 
fectant in  surgery,  for  cleansing  wounds  and  washing  the  hands 
and  surgical  instruments.  It  is  employed  only  in  dilute 
solutions.  The  solutions  have  corrosive  properties  and  a  sharp, 
very  disagreeable,  "  metallic  "  taste.  Mercuric  chloride  is  also 
employed  in  preserving  anatomical  specimens,  herbaria,  stuffed 
animals,  wood,  etc.  With  albumen  it  unites,  forming  insoluble 
compounds,  hence  the  use  of  white  of  egg  and  milk  in  case  of 


406  OUTLINES  OF   CHEMISTRY 

poisoning  with  corrosive  sublimate.  With  chlorides  of  the  alka- 
lies, mercuric  chloride  forms  double  salts,  like  HgCl2  •  KC1  •  H2O. 

Mercuric  bromide  HgBr2  is  similar  to  mercuric  chloride.  It 
is  isomorphous  with  the  latter  and  melts  at  325°. 

Mercuric  iodide  HgI2  may  be  prepared  by  the  direct  union  of 
iodine  and  mercury,  or  by  adding  potassium  iodide  to  a  solution 
of  mercuric  chloride,  when  a  yellowish  precipitate  forms  which 
soon  becomes  bright  red.  Though  insoluble  in  water,  it  readily 
dissolves  in  alcohol,  from  which  solutions  it  crystallizes  in  bright 
red,  tetragonal  pyramids.  In  aqueous  potassium  iodide  solu- 
tions, mercuric  chloride  is  readily  soluble,  forming  a  yellow 
solution  containing  the  double  salt  HgI2  •  2  KI.  This  solution 
may  be  concentrated  till  it  has  a  specific  gravity  of  over  3.  It 
is  known  as  Thoulet's  solution  and  is  used  by  mineralogists  in 
determining  the  specific  gravity  of  small  pieces  of  minerals. 
Potassium  mercuric  iodide  solution  to  which  caustic  potash 
has  been  added  is  Nessler's  reagent,  which  is  used  in  determin- 
ing ammonia  in  the  analysis  of  potable  waters.  With  ammonia 
such  solutions  give  a  brown  precipitate  having  the  composi- 


tion  O/  NH2  •  I.     If  the  ammonia  solutions  are  very  di- 


lute,  only  a  yellowish  brown  coloration  is  observed. 

Mercuric  Cyanide  Hg(CN)2,  the  only  cyanide  of  a  heavy 
metal  that  is  soluble  in  water,  is  obtained  by  the  action  of 
hydrocyanic  acid  on  mercuric  oxide.  It  forms  tetragonal 
prisms,  which  when  heated  yield  mercury  and  cyanogen. 

Nitrates  of  Mercury.  —  Mercurous  nitrate  HgNO3  •  H2O  forms 
monoclinic  crystals,  obtained  by  action  of  nitric  acid  upon  an 
excess  of  mercury  in  the  cold.  On  diluting  its  solutions  with 
water,  yellow  basic  salts  separate  out,  like  Hg(OH)HgNO3. 
This  hydrolysis  is  counteracted  by  keeping  a  slight  excess  of 
nitric  acid  in  the  solutions. 

Mercuric  nitrate  Hg(NO3)2  is  formed  by  dissolving  mercuric 
oxide  in  nitric  acid,  also  by  the  action  of  an  excess  of  hot,  con- 
centrated nitric  acid  on  mercury.  From  the  solutions,  deli- 
quescent crystals  2  Hg(NO3)2  -+•  H2O  may  be  obtained.  On 
dilution  with  water,  these  suffer  hydrolysis,  forming  a  series  of 
basic  salts  whose  composition  varies  according  to  the  rela- 
tive amount  of  water  present;  so  we  have,  for  instance, 


THE  METALS  OF  THE  MAGNESIUM  GROUP       407 

Hg(NO3)2  .  2  HgO  +  H2O.  On  boiling  with  much  water,  these 
basic  salts  are  decomposed,  yielding  nitric  acid  and  mercuric 
oxide.  Addition  of  nitric  acid,  of  course,  reverses  the  re- 
action. Under  the  name  Millon's  reagent,  mercuric  nitrate 
solution  is  used  in  testing  for  albumins,  which  are  coagulated 
by  it. 

Mercuric  Fulminate  HgC2O2N2,  also  called  fulminating  mer- 
cury, is  made  by  the  action  of  nitric  acid  on  mercury  in  pres- 
ence of  alcohol.  It  is  a  white  powder,  which  is  very  explosive 
when  dry.  It  is  used  in  making  percussion  caps.  The  explo- 
sion of  a  small  amount  of  the  substance  will  cause  gun  cotton 
or  nitroglycerine  to  explode. 

Sulphates  of  Mercury.  —  Mercurous  sulphate  Hg2SO4  is  formed 
by  treating  an  excess  of  mercury  with  sulphuric  acid,  or  by 
precipitating  a  mercurous  nitrate  solution  with  sulphuric  acid. 
The  colorless  crystals  so  obtained  are  sparingly  soluble  in 
water.  The  salt  is  used  in  making  standard  cells  for  compar- 
ison of  electromotive  forces. 

Mercuric  sulphate  HgSO4  is  made  by  digesting  mercury 
or  mercuric  oxide  with  an  excess  of  sulphuric  acid.  The  salt 
so  obtained  is  a  white,  crystalline  mass.  On  treatment  with 
water,  it  is  hydrolyzed,  forming  sulphuric  acid  and  a  yellow 
basic  salt  HgO  -  HgSO4,  known  as  Turpeth  mineral,  which  is 
insoluble.  With  sulphates  of  the  alkalies,  mercuric  sulphate 
forms  double  salts  like  K2SO4  -  HgSO4  6  H2O.  These  are 
isomorphous  with  the  analogous  magnesium  double  salts. 

Mercuric  Sulphide  HgS  is  very  stable.  It  is  readily  obtained 
by  triturating  mercury  and  moist  sulphur  together,  by  heating 
mercury  with  sulphur,  or  by  conducting  hydrogen  sulphide 
into  a  solution  of  a  mercury  salt.  Mercurous  sulphide  is  not 
known  with  certainty;  the  black  precipitate  produced  when  a 
mercurous  salt  is  treated  with  hydrogen  sulphide  or  an  alkaline 
sulphide  is  mercuric  sulphide  mixed  with  mercury.  When 
made  by  precipitation,  mercuric  sulphide  is  a  black,  amorphous 
powder,  which  is  insoluble  even  in  concentrated  acids  on  boil- 
ing. It  is  soluble  in  aqua  regia,  however.  Black  mercuric  sul- 
phide may  be  sublimed,  and  thus  converted  into  dark  red 
rhombohedral  crystals  that  are  identical  with  cinnabar,  which 
occurs  in  nature.  Towards  acids  the  red  and  black  varieties 
act  alike.  Red  sulphide  of  mercury  serves  as  a  pigment  in 


408  OUTLINES   OF   CHEMISTRY 

paints  under  the  name  of  vermilion.  It  has  been  used  thus 
since  ancient  times. 

Compounds  of  Mercury  Salts  with  Ammonia.  —  When  mercury 
salts  are  treated  with  ammonia,  compounds  are  formed  which 
may  be  regarded  as  ammonium  salts  in  which  one  or  more  of 
the  hydrogen  atoms  are  replaced  by  mercury.  So  when  mer- 
curous  chloride  is  treated  with  ammonia  water,  a  black 
insoluble  powder,  mercurous  ammonium  chloride,  is  formed :  — 

2  HgCl  +  2  NH3  =  (NH2Hga)  Cl  +  NH4C1. 
Similarly,  mercurous  nitrate  yields  black  mercurous  ammonium 
nitrate  (NH2Hg2)NO3  plus  mercury.     n 

Mercuric  ammonium  chloride  (NH2Hg)Cl,  known  as  infu- 
sible white  precipitate,  is  formed  by  adding  ammonia  to  mer- 
curic chloride  solution:  — 

HgCl2  +  2  NH8  =  (NH2Hg)Cl  +  NH4C1. 
The  so-called  fusible  white  precipitate,    mercuric  diammonium 

chloride  (NH3Cl)2Hg,  is  obtained  by  treating  a  boiling  hot 
ammoniacal  solution  of  ammonium  chloride  with  mercuric 
chloride. 

Physiological  Properties  of  Mercury  Compounds.  —  It  has 
already  been  stated  that  mercury  compounds  are  poisonous. 
When  mercury  is  introduced  into  the  system,  it  produces  a 
peculiar  taste  in  the  mouth  commonly  described  as  metallic. 
Salivation  follows,  and  the  gums,  teeth,  liver,  kidneys,  and  other 
organs  frequently  also  become  involved.  In  small  doses  mer- 
cury compounds  stimulate  the  action  of  various  glands ;  this 
furnishes  the  basis  for  the  internal  use  of  calomel.  The  intro- 
duction of  mercury  compounds  into  medicine  dates  back  to 
Paracelsus,  who  lived  in  the  first  half  of  the  sixteenth  century. 

Tests  for  Mercury.  —  Mercury  is  very  readily  detected  in  its 
compounds.  By  mixing  any  mercury  compound  with  soda  and 
heating,  mercury  is  driven  off,  which  condenses  in  drops  in  the 
cooler  parts  of  the  ignition  tube.  With  iodine  these  drops 
form  red  iodide  of  mercury. 

In  solutions  of  mercurous  salts,  chlorides  produce  a  white 
precipitate  of  calomel  which  turns  black  on  treatment  with 
ammonia,  as  already  stated. 

The  fact  that  stannous  chloride  reduces  mercuric  chloride  to 


THE  METALS  OF  THE  MAGNESIUM  GROUP       409 

mercurous  chloride,  and  that  the  latter  is  further  reduced  to 
mercury  by  adding  more  stannous  chloride,  is  frequently  used 
in  testing  for  mercury.  Further,  bright  copper  when  intro- 
duced into  a  solution  of  a  mercury  salt  becomes  coated  with 
mercury.  The  sulphide  of  mercury  is  also  very  characteristic, 
as  already  mentioned. 

General  Remarks.  —  It  will  be  observed  that  the  compounds 
of  magnesium  and  zinc  bear  close  resemblances  to  each  other, 
and  also  a  fair  resemblance  to  the  compounds  of  cadmium.  On 
the  other  hand,  the  mercury  compounds  do  not  resemble  those 
of  zinc,  magnesium,  and  cadmium.  To  be  sure,  in  the  mercuric 
compounds,  mercury  is  bivalent,  and  between  this  series  and  the 
compounds  of  the  other  metals  mentioned,  some  analogies  are 
apparent.  It  will  be  seen  later  that  the  compounds  of  mercury 
bear  a  closer  resemblance  to  those  of  copper.  Glucinum  stands 
rather  isolated,  being,  as  has  already  been  mentioned,  a  transi- 
tion element  between  this  group  and  that  of  the  earth  metals. 
Thus  it  is  evident  that  the  metals  of  the  magnesium  group  do 
not  form  as  closely  related  a  family  as  some  of  the  others  that 
have  already  been  studied. 


CHAPTER   XXIV 

SOLUTIONS,   ELECTROLYSIS,   AND   ELECTRO-CHEMICAL 

THEORIES 

Nature  and  Kinds  of  Solutions.  —  As  stated  in  Chapter  I, 
solutions  were  formerly  regarded  as  chemical  compounds  accord- 
ing to  variable  proportions.  This  designation  expresses  the 
fact  that  the  process  of  solution  is  accompanied  by  all  the  phe- 
nomena that  are  observed  when  chemical  change  takes  place, 
and  that  the  composition  of  the  final  homogeneous  product 
obtained  may  gradually  be  varied  at  will  within  certain  limits, 
as  has  already  been  explained.  The  term  solution  has  of  recent 
years  received  a  somewhat  broader  meaning  than  it  formerly 
had.  So  mixtures  of  gases  are  sometimes  spoken  of  as  solu- 
tions. Glass,  various  alloys,  isomorphous  mixtures  of  crystal- 
line substances,  gases  absorbed  by  solids,  etc.,  are  frequently 
termed  solid  solutions.  A  solution  is  commonly  defined  as  a 
homogeneous  mixture,  the  sconstituent  parts  of  which  cannot 
be  separated  by  mechanical  means.  The  term  mixture  as  here 
used  is,  however,  not  to  be  confused  with  a  mere  mechanical 
mixture  like  that  of  sulphur  and  iron  filings  ground  together, 
as  already  remarked  in  Chapter  I.  The  word  mixture,  as 
used  in  connection  with  a  solution,  indicates  solely  that  the 
proportions  of  the  ingredients  may  be  varied  arbitrarily,  at 
least  to  some  extent. 

Any  gas  will  mix  with  any  other  gas  or  mixture  of  gases  in  all 
proportions,  as  would  naturally  be  expected,  for  in  a  gas  the 
molecules  are  relatively  remote  from  one  another.  But  gases 
will  not  be  absorbed  by  liquids  or  solids  in  all  proportions.  Here 
the  specific  nature  of  the  gas  and  that  of  the  liquid  or  solid  under 
consideration  is  of  prime  importance.  The  specific  nature  of 
liquids  and  solids  is  also  the  determining  factor  in  fixing  the  solu- 
bility of  liquids  or  solids  in  other  liquids  or  solids.  Moreover, 
in  all  cases  of  solution,  the  temperature  is  very  important,  being 
second  only  to  the  effect  of  the  nature  of  the  substances.  Pressure 
is  also  a  factor,  which  is  particularly  important  when  gases 

410 


SOLUTIONS  AND   ELECTRO-CHEMICAL   THEORIES          411 

come  into  play,  for  these  are  highly  compressible.  On  the  other 
hand,  the  pressure  factor  is  of  less  importance  when  liquids  and 
solids  only  are  used.  Indeed,  the  factors  mentioned  are  of  the  same 
relative  importance  in  the  process  of  solution  as  in  chemical  action. 
Absorption  of  Gases  by  Liquids. — The  amount  of  a  gas  absorbed 
by  a  liquid  increases  as  the  temperature  is  lowered  and  as  the 
pressure  is  increased.  At  constant  temperature,  the  amount  of 
gas  absorbed  by  a  liquid  is  directly  proportional  to  the  pressure. 
This  law  was  discovered  in  1803  by  Henry,  whose  name  it  bears. 
Henry's  law  holds  only  in  case  the  gas  is  not  very  soluble  in  the 
liquid,  like  oxygen  or  hydrogen  in  water,  or  nitrogen  in  alco- 
hol. When  gases  are  copiously  soluble,  like  hydrochloric  acid 
or  ammonia  in  water,  the  law  does  not  hold.  Gases  that  follow 
Henry's  law  have  a  very  low  heat  of  solution,  showing  that  but 
little  affinity  exists  between  the  gas  and  the  liquid.  Gases 
that  do  not  follow  the  law  have  a  high  heat  of  solution,  indicat- 
ing the  prominence  of  the  affinity  factor  ;  for  the  heat  developed 
is  frequently  greater  than  that  produced  when  the  dry  gas  is 
liquefied  by  pressure.  Thus  we  have  :  — 

NH3  +  (aq)  =  (NH3aq)  +  8.8  Cal., 

whereas  the  heat  of  liquefaction  of  17  grams  of  ammonia  is  only 
4.4  Cal. 

Solutions  of  Liquids  in  Liquids.  —  Many  liquids  mix  perfectly 
with  one  another  in  all  proportions,  forming  solutions.  Thus 
water  will  mix  with  glycerine,  alcohol,  or  acetone ;  ether  with 
kerosene  or  fatty  oils  ;  olive  oil  with  linseed  oil ;  carbon  disul- 
phide  with  ether,  hydrocarbon  oils,  or  fats.  In  general,  liquids 
like  hydrocarbon  oils  and  their  halogen  substitution  products, 
ethers,  esters,  carbon  disulphide,  and  the  various  fatty  arid  oily 
products  of  plant  or  animal  origin,  are  readily  soluble  in  one 
another  in  all  proportions.  Liquids  that  are  miscible  in  all  pro- 
portions are  called  consolute  liquids. 

Now,  hydrocarbon  oils  and  animal  and  vegetable  oils,  carbon 
disulphide,  and  carbon  tetrachloride  are  practically  insoluble  in 
water.  When  such  substances  are  added  to  water,  the  liquids 
separate  in  two  distinct  layers.  When  the  specific  gravities  of 
the  two  non-miscible  liquids  are  nearly  alike  arid  the  liquids  are 
shaken  together,  a  mixture  of  milky  appearance  is  obtained, 
which  on  closer  inspection  is  found  to  consist  of  small  globules 


412  OUTLINES   OF   CHEMISTRY 

of  one  liquid  suspended  in  the  other.  This  is  an  emulsion, 
The  greater  the  difference  in  specific  gravity  between  the  two 
liquids,  the  sooner  will  they  again  separate  into  two  layers.  The 
specific  nature  of  the  liquids  also  has  to  do  with  the  length  of 
time  the  globules  remain  in  suspension,  for  the  greater  the  ad- 
hesion between  the  liquids,  the  longer  the  emulsion  lasts.  In 
milk,  the  butter  fat  is  present  in  the  form  of  minute  globules  in 
suspension,  that  is,  in  emulsified  condition.  On  standing,  the 
fat  gradually  comes  to  the  top,  forming  a  layer  of  cream.  By 
centrifugal  force,  the  constituents  of  an  emulsion  can  be  rapidly 
separated,  which  fact  rs  used  in  separating  butter  fat  from  milk 
by  means  of  the  cream  separators  in  use  in  creameries. 

A  few  liquids  like  alcohol  and  acetone  are  consolute  either  with 
hydrocarbons  and  fats  or  with  water.  It  is  to  be  observed  that, 
when  alcohol  or  acetone  is  added  to  water,  the  resulting  solution 
will  then  take  up  much  more  hydrocarbon  or  fatty  oil  than  will 
pure  water.  A  knowledge  of  this  fact  is  often  of  practical 
value. 

Ether  is  soluble  in  water  to  a  limited  extent.  When  an  excess 
of  ether  is  added  to  water,  two  layers  form,  the  upper  one  finally 
becoming  a  saturated  solution  of  water  in  ether  and  the  lower 
one  a  saturated  solution  of  ether  in  water.  The  solubility  in 
this  case,  as  in  all  others,  is  a  function  of  the  temperature. 
Phenol  C6H5OH  and  water  are  also  only  partially  miscible  at 
ordinary  temperatures;  however,  at  and  above  68.4°  the  two 
liquids  become  perfectly  consolute.  Again,  above  20°  triethyl- 
amine  N(C2H5)3  and  water  act  like  ether  and  water,  that  is,  they 
are  not  consolute.  But  below  20°  triethylamine  and  water  are 
miscible  in  all  proportions.  In  this  case  the  miscibility  increases 
as  the  temperature  is  lowered,  though  more  frequently  the  op- 
posite is  true  —  as  in  the  pair,  phenol  and  water. 

Solutions  of  Solids  in  Liquids.  —  Whether  a  given  solid  will  dis- 
solve in  a  given  liquid  and  to  what  extent,  depends  primarily  upon 
the  specific  nature  of  the  two  substances,  and  also  upon  the  tempera- 
ture. The  pressure  is  also  a  factor,  strictly  speaking,  but  ordi- 
narily it  is  of  little  practical  importance.  A  liquid  that  dissolves 
a  solid  adheres  to  or  wets  the  latter.  Nevertheless,  in  many 
cases  a  liquid  will  wet  a  solid  without  dissolving  it  appreciably. 
We  may  look  upon  adhesion  as  an  unsuccessful  attempt  to  form 
a  solution,  for  here  the  attraction  between  solid  and  liquid  is 


SOLUTIONS   AND   ELECTRO-CHEMICAL   THEORIES 


413 


insufficient  to  overcome  their  cohesions  and  thus  form  a  solu- 
tion. 

Many  solids  have  definite  solubility  in  a  given  liquid,  which 
varies  with  the  temperature.     Thus  the  ordinary  salts  and  other 


140 


20 


40  60 

Temperature  in  degrees  C. 
FIG.  138. 


80 


100 


crystalline  solids,  like  sugar  or  urea,  have  a  definite  solubility  in 
water.     On  the  other  hand,  substances  like  gelatine,  gum  arabic, 


414  OUTLINES   OF  CHEMISTRY 

and  many  others  of  similar  non-crystalline  character  do  not 
have  a  definite  solubility ;  they  are  practically  consolute  with 
water.  So  when  water  is  dropped  on  glue,  the  latter  swells  up, 
and  as  more  and  more  water  is  gradually  added  all  stages  of 
plasticity  are  obtained,  till  a  thick  sirup,  and  finally  a  limpid 
liquid,  results.  In  general,  crystalline  substances  have  sharply 
defined  melting  points  as  well  as  definite  solubilities,  whereas 
non-crystalline  substances  frequently  lack  definite  melting 
points  and  solubilities,  though  this  rule  is  not  without  its 
exceptions. 

The  solubilities  of  solids  in  liquids  vary  very  greatly  with  the 
nature  of  the  substances  brought  together,  as  well  as  with  the  tem- 
perature. Some  substances,  like  common  salt  in  water,  are 
about  as  soluble  at  lower  as  at  higher  temperatures. 

More  frequently,  the  solubility  increases  with  rise  of  temper- 
ature, as  in  the  case  of  saltpeter  in  water. 

The  number  of  grams  of  a  substance  dissolved  by  100  grams 
of  a  given  liquid  at  a  certain  temperature  is  called  the  solubility 
of  that  substance  at  the  given  temperature.  Solubility  curves 
are  commonly  constructed  by  charting  temperatures  as  abscissas 
and  solubilities  as  ordinates.  Figure  138  shows  a  number  of 
solubility  curves  of  well-kru^wn  salts.  It  will  be  observed  that 
these  by  no  means  have  the  same  trend.  Even  much  greater 
varieties  than  these  are  met  frequently. 

Degrees  of  Saturation.  —  In  determining  the  solubility  of  a 
solid,  the  usual  practice  is  to  shake  the  liquid  with  an  excess  of 
the  solid  till  no  more  is  dissolved.  This  process  is  conducted 
at  the  temperature  at  which  the  solubility  is  sought.  The 
liquid  finally  obtained  is  said  to  be  a  saturated  solution  at  that 
temperature.  The  strength  of  the  saturated  solution  is  found 
by  analyzing  a  given  weight  of  it.  This  frequently  consists 
simply  of  evaporating  the  solution  to  dry  ness  and  weighing  the 
residue.  If  the  saturated  solution  thus  obtained  at  a  given 
temperature,  say  30°,  is  carefully  decanted  from  the  excess  of 
solid  substance  so  that  none  of  the  latter  is  present,  and  then 
this  liquid  is  heated  above  30°,  the  solution  is  said  to  be 
unsaturated;  for  at  this  higher  temperature  it  would  take  up 
more  of  the  solid  if  some  of  the  latter  were  introduced.  On 
the  other  hand,  if  the  saturated  solution  is  cooled  slightly 
below  its  temperature  of  saturation  out  of  contact  with  the  solid 


SOLUTIONS  AND  ELECTRO-CHEMICAL  THEOEIES          415 

it  contains,  the  liquid  usually  remains  clear.  The  solution  is 
now  said  to  be  supersaturated,  for  when  brought  into  contact 
with  some  of  the  solid  (even  traces  will  suffice),  more  of  the 
latter  will  drop  out  and  the  solution  will  then  become  saturated 
at  this  lower  temperature.  This  is  another  method  of  making 
a  saturated  solution.  The  degree  of  supersaturation  that  may 
be  obtained  by  cooling  saturated  solutions  as  described  varies 
in  the  case  of  different  solutions.  The  phenomena  of  supersatu- 
rated solutions  are  analogous  to  those  of  the  supercooling  of  liquids. 
So,  out  of  contact  with  ice,  water  may  be  cooled  several  degrees 
below  zero  and  still  be  a  liquid.  On  touching  this  supercooled 
water  with  a  piece  of  ice,  however,  the  whole  congeals. 

Unsaturated,  saturated,  and  supersaturated  solutions  are 
similarly  obtained  in  case  of  solutions  of  gases  in  liquids  and  of 
liquids  in  liquids. 

Solid  Solutions.  — These  have  already  been  mentioned.  They 
are  commonly  prepared  by  melting  two  solids  together  and  al- 
lowing the  liquid  to  congeal.  Sometimes  by  crystallizing  cer- 
tain substances  from  solution,  crystals  of  variable  composition, 
solid  solutions,  result.  Again,  gases  are  absorbed  by  solids  to 
form  so-called  solid  solutions.  In  these  cases  we  may  also  have 
either  restricted  solubility,  or  miscibility  in  all  proportions. 

Precipitation.  —  The  process  of  precipitation  is  the  opposite  to 
that  of  solution.  Thus  when  a  salt  has  been  dissolved  in  water, 
the  latter  may  be  evaporated  from  the  solution  till  the  salt 
separates  out.  Again,  common  salt  is  soluble  in  water,  but  not 
in  alcohol;  hence  by  adding  alcohol  to  salt  brine,  some  of  the 
salt  is  precipitated.  Similarly,  camphor  is  soluble  in  alcohol  but 
not  in  water ;  consequently  camphor  is  precipitated  by  adding 
water  to  a  solution  of  camphor  in  alcohol.  Sodium  chloride  is 
insoluble  in  liquid  hydrochloric  acid,  hence  it  is  readily  compre- 
hended why  the  salt  is  precipitated  when  hydrochloric  gas  or 
even  a  concentrated  solution  of  it  is  added  to  salt  brine.  Cal- 
cium phosphate  is  soluble  in  nitric  or  hydrochloric  acids  but  not 
in  water;  hence  by  neutralizing  such  an  acid  solution  of  calcium 
phosphate  with  ammonia  or  caustic  soda,  calcium  phosphate  is 
precipitated.  These  are  all  simple,  typical  illustrations. 

Precipitation  may,  however,  also  be  caused  ~by  double  decom- 
position in  solutions.  So  silver  chloride  is  precipitated  thus  :  — 

AgNO3  +  NaCl  =  NaNO3  +  AgCl. 


416  OUTLINES  OF  CHEMISTRY 

This  reaction  occurs  because  under  the  conditions  that  obtain 
silver  chloride  and  the  solution  of  sodium  nitrate  are  the  stabler 
products.  The  precipitation  does  not  take  place  simply  because 
there  is  an  opportunity  for  the  insoluble  silver  chloride  to  form, 
it  proceeds  also  because  there  is  the  chance  to  form  the  stable 
and  soluble  sodium  nitrate.  In  other  words,  reactions  of  this 
kind  must  be  considered  as  a  whole.  When  silver  nitrate  solu- 
tion is  shaken  with  carbon  tetrachloride,  silver  chloride  does 
not  form.  Clearly  there  is  plenty  of  silver  and  chlorine  to 
form  silver  chloride,  but  the  other  product  would  have  to  be  a 
nitrate  of  carbon,  which  is  unknown,  probably  because  of  its 
great  instability.  This  lack  of  action  between  silver  nitrate  and 
carbon  tetrachloride  is  not  to  be  ascribed  to  the  fact  that  the  latter 
compound  is  a  non- electrolyte,  as  is  often  done,  for  similar  changes 
do  occur  in  the  best  of  insulators,  as  will  be  shown  below.  By 
heating  carbon  tetrachloride  with  silver  nitrate  and  nitric  acid 
to  high  temperatures  in  a  sealed  tube  (method  of  Carius)  silver 
chloride  is  indeed  obtained,  for  here  there  are  conditions  under 
which  carbon  is  oxidized.  Thus  carbon  tetrachloride  is  de- 
stroyed, carbon  dioxide,  oxides  of  nitrogen,  and  silver  chloride 
resulting  as  the  stable  products  under  the  conditions  to  which 
the  mixture  was  subjected.  The  method  of  precipitation  by 
double  decomposition  in  solutions  is  very  commonly  employed  in 
chemistry.  In  all  cases  the  explanation  is  similar  to  that  of  the 
precipitation  of  silver  chloride. 

Liquids  also  may  be  thrown  out  of  solution.  So,  many  oils  are 
insoluble  in  water  but  soluble  in  alcohol.  By  adding  water  to 
such  alcoholic  solutions,  the  oil  separates  out  in  the  form  of  a 
layer. 

G-ases  also  may  be  liberated  from  their  solutions  in  liquids.  In 
many  cases  this  can  be  done  by  simply  heating  the  solution. 
But  it  may  also  be  accomplished  by  adding  some  other  appro- 
priate liquid,  solid,  or  gas  to  the  solution.  So,  for  instance,  by 
adding  concentrated  sulphuric  acid  to  a  strong  solution  of 
hydrochloric  acid,  hydrochloric  acid  gas  is  liberated ;  similarly, 
ammonia  is  set  free  when  concentrated  caustic  potash  solution 
is  added  to  a  strong  aqueous  solution  of  ammonia. 

Colloidal  Solutions.  —  These  may  be  made  by  dialysis,  as  al- 
ready explained.  When  an  electric  arc  is  formed  between  two 
rods  of  silver  dipping  into  water  (Fig.  139)  a  dark-colored  liquid 


SOLUTIONS   AND   ELECTRO-CHEMICAL  THEORIES          417 

results  which  is  an  extremely  finely  divided  suspension  of  sil- 
ver. This  is  sometimes  called  a  colloidal  solution  of  silver. 
Similar  colloidal  solutions  of  other 
metals,  like  platinum  or  gold,  can 
be  formed. 

In  general,  colloidal  solutions 
boil  at  the  same  temperature  as  the 
pure  solvent.  They  vary  from 
mere  suspensions  to  true  solu- 

T        V      •  ^     '  11      -J     1  FlQ-    139' 

tions.     In  living  beings,  colloidal 

solutions  play  an  important  role.     Many  colloidal  solutions  are 

precipitated  by  slight  additions  of  various  salts,  sugar,  alcohol, 

etc.     The   statement   sometimes  made    that   only  electrolytes 

precipitate   colloidal    solutions    is    erroneous.       For   instance, 

sugar  added  to  colloidal  solutions  of  ferric  hydroxide  causes 

precipitation. 

Boiling  Points  of  Solutions.  —  The  boiling  point  of  a  liquid  is 
the  temperature  at  which  its  vapor  tension  just  overcomes  atmos- 
pheric pressure.  It  consequently  varies  with  the  pressure.  At 
constant  pressure,  the  boiling  point  of  a  liquid  is  changed  by 
dissolving  substances  in  it.  Thus  solutions  are  obtained  which 
may  have  a  higher  or  lower  boiling  point  than  the  solvent, 
depending  on  the  nature  of  the  substances  brought  together. 
So  water  boiling  at  100°  and  alcohol  boiling  at  78°  yield  solu- 
tions that  boil  between  these  temperatures.  Formic  acid  boil- 
ing at  101°  may  yield  an  aqueous  solution  that  boils  at  107°. 
Pyridine  boiling  at  115°  and  water  may  yield  a  solution  boiling 
at  93°. 

Use  of  Boiling  Points  of  Solutions  in  Molecular  Weight  Determi- 
nations. —  When  a  substance  that  is  practically  non-volatile,  as 
compared  with  the  solvent  liquid,  is  dissolved  in  the  latter,  the 
solution  so  obtained  boils  higher  than  the  pure  solvent.  So, 
when  sugar  is  added  to  water,  the  solution  obtained  boils  above 
100°.  The  rise  of  the  boiling  point  is  proportional  to  the  amount 
of  solute  added.  Furthermore,  the  elevations  of  the  boiling  point 
are  approximately  the  same  for  equimolecular  quantities  of  solute 
each  dissolved  in  an  equal  quantity  of  solvent.  This  offers  a 
simple  method  of  determining  the  molecular  weights  of  sub- 
stances in  solution.  So  342  grams  of  sugar  C12H22On,  60  grams 
of  urea  CO(NH2)2,  94  grams  of  glycerine  C3H5(OH)8,  120 

2E 


418  OUTLINES   OF   CHEMISTKY 

grams  of  magnesium  sulphate,  159  grams  of  copper  sulphate 
CuSO4,  each  dissolved  in  5000  grams  of  water,  will  yield  solu- 
tions that  boil  at  100.104°.  This  rise  of  0.104°  in  5000  grams 
of  water  would  be  proportional  to  a  rise  of  50  x  0.804°  or  5.20° 
in  100  grams  of  water.  The  value  5.20°  is  termed  the  molecular 
rise  of  the  boiling  point  when  water  is  used  as  the  solvent.  For 
other  solvents  this  value  is  quite  different.  For  instance,  for 
ether  it  is  21.6°,  for  alcohol  11.7°,  for  carbon  tetrachloride  41.0°, 
for  benzene  26.7°.  It  is  possible  to  compute  these  values  from 
the  absolute  boiling  point  T  and  the  latent  heat  of  evaporation 
L  of  the  solvent  by  means  of  the  formula, 

2  Tz 

Molecular  rise  of  boiling  point  =    AA  T  • 

1UU  lj 

The  theoretical  considerations  upon  which  this  formula  is  based 
will  not  be  entered  into  here. 

To  ascertain  the  molecular  weight  of  any  substance  soluble 
in  water,  dissolve  so  much  of  the  substance  in  5000  grams  of 
water  till  the  solution  boils  at  100.104°.  The  amount  of  solute 
added  is  its  molecular  weight  in  grams.  In  practice  one  would, 
of  course,  take  a  much  smaller,  convenient  amount  of  water 
and  solute  to  determine  the  boiling  point  of  the  solution,  and 
compute  the  quantity  of  solute  required  to  elevate  the  boiling 
point  5.2°  per  every  100  grams  of  water.  Instead  of  the  rise 
of  the  boiling  point  of  a  solution,  the  lowering  of  the  vapor 
tension  caused  by  the  introduction  of  the  solute  may  be  used  to 
determine  the  molecular  weight  of  the  solute. 

The  Freezing  Point  of  Solutions.  —  The  freezing  point  of  a 
solution  is  lower  than  that  of  the  pure  solvent.  The  lowering 
of  the  freezing  point  is  also  about  the  same  for  equimolecular 
quantities  of  solutes  each  dissolved  in  an  equal  amount  of  solvent. 
Thus,  in  aqueous  solutions,  the  molecular  lowering  of  the  freez- 
ing point  is  about  18.9°  per  100  grams  of  water.  The  value 
varies  for  different  solvents,  but  may  be  computed  approximately 
from  the  absolute  freezing  point  and  the  latent  heat  of  fusion 
of  the  solvent  according  to  theoretical  considerations  of  van't 
Hoff,  which  will  not  be  reproduced  here,  the  formula  being  — 

2  T  2 

Molecular  lowering  of  the  freezing  point  =         *    . 

100  jj-, 

The  method  of  determining  molecular  weights  in  solution  from 


SOLUTIONS  AND  ELECTRO-CHEMICAL  THEORIES 


419 


the   lowering  of  the  freezing  point  is  analogous   to   that   of 
making  such  estimations  from  the  rise  of  the  boiling  point. 

Discussion  of  Molecular  Weights  Determined  in  Solutions.  — 
The  freezing  and  boiling  point  methods  yield  only  approxi- 
mate results  at  best,  and  can  be  used  with  success  only  in 
dilute  solutions.  Special  care  must  be  exercised  in  interpret- 
ing the  results  obtained,  for  they  are  often  anomalous.  So, 
many  substances  like  gelatine,  dialyzed  silicic  acid,  or  soap 
when  dissolved  in  water  hardly  cause  any  rise  of  the  boiling 
point.  These  substances  consequently  have  enormous  molecu- 
lar weights  as  found  by  these  methods.  Such  substances  are 
consequently  frequently  regarded  as  highly  associated  or  poly- 
merized. Again,  many  substances,  among  them  a  large  number 
of  the  ordinary  acids,  salts,  and  bases,  when  dissolved  in  water 
and  tested  by  the  freezing  or  boiling  point  method,  yield  much 
lower  molecular  weights  than  the  formulae  usually  ascribed  to 
them  indicate.  These  substances  are  by  some  chemists  re- 
garded as  dissociated  in  solution  ;  they  seek  to  connect  the 
abnormally  low  molecular  weights  observed  with  the  fact  that 
such  solutions  are  frequently  electrolytic  conductors.  So,  for 
instance,  common  salt,  saltpeter, 
caustic  potash,  nitric  acid,  etc., 
yield  molecular  weights  that  are 
usually  somewhat  over  half  of 
those  represented  by  the  formulae 
of  these  compounds  (see  the  theory 
of  Arrhenius). 

Osmosis  and  Osmotic  Pressure. 
—  If  carbon  disulphide  C  (Fig. 
140)  be  placed  in  a  tube  and  a 
layer  of  water  B  be  poured  upon 
it,  and  again  a  layer  of  ether  A  be 
carefully  poured  upon  the  water, 
there  will,  after  long  standing, 
eventually  be  but  two  layers  A1 
and  B'  (Fig.  141).  The  lower 
layer  Bf  consists  of  carbon  disul- 
phide and  ether  saturated  with 

water,  whereas  the  upper  layer  A1  consists  of  water  saturated 
with  carbon  disulphide   and   ether.      The  change  which  has 


.Fia.  140. 


FIG.  141. 


420  OUTLINES   OF   CHEMISTRY 

occurred  is  easily  explained.  Ether  dissolves  very  readily 
in  carbon  disulphide,  but  much  less  readily  in  water.  Again, 
carbon  disulphide  and  water  hardly  dissolve  each  other  at  all. 
In  Fig.  140  the  water  layer  B  dissolves  ether  and  in  turn  the 
ethereal  layer  A  also  takes  up  some  water.  When  the  ether  has 
gone  into  B  till  it  touches 'the  carbon  disulphide  (7,  the  latter 
extracts  ether  from  the  aqueous  layer  B.  Thus  the  carbon 
disulphide  layer  O  becomes  enriched  with  ether,  whereas  the 
aqueous  layer  becomes  depleted  in  ether.  The  depletion  is 
made  good  by  a  continuous  supply  of  ether  from  A  till  the 
latter  is  exhausted  and  equilibrium  is  produced.  Water 
charged  with  ether  dissolves  carbon  disulphide  to  a  greater 
extent  than  pure  water  does,  and  so  the  aqueous  layer  B 
always  contains  some  carbon  disulphide.  Moreover,  before 
equilibrium  is  reached,  some  of  this  carbon  disulphide  is  ex- 
tracted from  the  layer  B  by  the  ethereal  layer  A.  Thus,  in  the 
attempt  to  reach  equilibrium,  ether  is  passing  from  A  through 
B  into  <7,  and  on  the  other  hand  carbon  disulphide  is  traveling 
from  0  through  B  into  A.  The  former  current  is  much  the 
stronger,  and  so  the  layer  A  gradually  disappears. 

The  example  just  cited  is  a  typical  case  of  osmosis.  The  layer 
B  is  the  septum  which  separates  the  liquids  A  and  O.  Whether 
the  latter  substances  will  pass  through  B  or  not  is  determined 
by  the  specific  nature  of  B  and  also  that  of  0  and  of  A.  Further- 
more, the  specific  nature  of  the  septum  B  and  that  of  the  liquids 
A  and  (7  which  bathe  it  also  determines  the  direction  which 
the  major  current  will  take.  In  all  osmotic  processes  there  is  a 
major  and  a  minor  current,  but  in  some  cases  the  latter  is  so  slight 
as  to  be  almost  negligible.  Under  these  conditions  the  osmotic 
process  appears  to  be  one-sided  and  the  septum  used  is  said  to  be 
semipermeable. 

It  is  clear  that  if  the  layer  B  in  Fig.  140  could  be  held  rig- 
idly in  place,  pressure  would  result  on  the  walls  of  the  compart- 
ment O  because  of  the  influx  of  ether.  Figure  142  shows  an 
apparatus  of  glass  containing  mercury,  carbon  disulphide  in  (7, 
and  ether  in  A,  while  B  is  a  slice  of  cork  which  is  saturated  with 
water,  and  tightly  jammed  into  the  tube.  As  the  apparatus 
stands,  the  mercury  rises  in  the  limb  at  the  right  because  the 
major  current  consists  of  the  passage  of  ether  from  A  through 
B  into  (7,  as  already  explained.  The  pressure  so  produced  is 


SOLUTIONS  AND   ELECTRO-CHEMICAL   THEORIES 


421 


called  the  osmotic  pressure.     It  is  clear  that  this  pressure  is  the 

result   of  differences   in   solubility.     The    apparatus   shown   in 

Figure  142   serves  only  to   demonstrate 

the  existence  of  osmotic  pressure,  not  to 

measure  its  value,  for  the  water-soaked 

cork  will  give  way  long  before  the  niaxi- 

mum  pressure  is  reached. 

Osmosis  has  been  studied  for  over  a 

hundred  years.     The  water,  ether,  carbon 

disulphide  experiment  was  performed  by 

L'Hermite  in  1854.     The  importance  of 

osmotic   investigations    in    physiological 

processes  was  recognized  very  early,  and 

physiologists  have  contributed  largely  to 

our  knowledge  of  osmosis.      When  septa, 

like  parchment  paper,  or  pieces  of  animal 

bladder,  are  used  to  separate  an  aqueous 

sugar   solution    S  from    pure   water    W 

(Fig.  143),  the  main  osmotic  current  is 

from  the  water 
through  the 
membrane  to  the 
solution,  for  the 
latter  extracts 

water  from  the  water- soaked  septum. 
At  the  same  time  much  sugar  passes 
through  the  septum  into  the  water  W. 
If,  however,  a  precipitate  of  copper  fer- 
rocyanide  Cu2Fe(CN)6-;r  H2O  is  formed 
in  the  pores  of  the  septum,  almost  no 
sugar  passes  from  the  cell  into  the 
outer  liquid  W\  that  is  to  say,  copper 
ferrocyanide  is  a  semipermeable  mem- 
brane, for  it  allows  water  to  pass 
through  it,  whereas  it  is  nearly  im- 
pervious to  sugar.  The  precipitate  of 
copper  ferrocyanide  is  formed  in  the 
membrane  by  first  placing  about  a 

three  per  cent  potassium  ferrocyanide  solution  in  the  cell  and 

then  immersing  the  latter  for  a  time  in  a  solution  of  copper 


FIG.  142. 


FIG.  143. 


422 


OUTLINES   OF   CHEMISTRY 


sulphate  of  about  equal  strength.  Thus  the  precipitate  forms 
in  the  pores  of  the  septum  :  — 

K4Fe(CN)6  +  2  CuSO4  =  2  K2SO4  +  Cu2Fe(CN)6. 

By  forming  this  precipitate  in  the  pores  of  a  small  unglazed 
porcelain  cup  C  (Fig.  144),  the  plant  physiologist,  Pfeffer,  in 
1877,  measured  the  maximum  osmotic  pressure  of  dilute  cane 

sugar  solutions  of  several 
concentrations  at  a  number 
of  different  temperatures. 
Figure  144  shows  the  cell, 
with  manometer  M attached, 
immersed  in  the  large  glass 
dish  filled  with  water.  T 
and  T  are  thermometers. 
In  1887  van't  Hoff  showed 
that  Pfeffer's  results  indi- 
cate that  the  osmotic  pres- 
sure is  proportional  to  the 
absolute  temperature,  and 
that  the  osmotic  pressure 
of  the  sugar  solution  is  the 
same  as  the  gas  pressure 
that  would  be  produced  if 
the  sugar  were  in  the  gase- 
ous state  and  confined  in 
the  same  volume  that  the 
solution  occupies.  In  other 
words,  van't  Hoff  showed 
that  when  dilute  aqueous 
cane  sugar  solutions  are 
separated  from  water  by 
means  of  copper  ferrocyanide 
membranes,  the  osmotic  pressures  developed  may  be  represented 
by  the  gas  equation  PV—RT^  where  P  is  the  osmotic  pres- 
sure measured,  V  the  volume  of  the  solution,  T  the  absolute 
temperature,  and  R  the  usual  gas  constant.  More  recently, 
Morse,  also  using  copper  ferrocyanide  membranes,  showed  by 
similar,  though  far  more  elaborate,  experiments  that  aqueous 
cane  sugar  solutions  do  develop  osmotic  pressures  that  approxi- 


FIG.  144. 


SOLUTIONS  AND   ELECTRO-CHEMICAL  THEORIES          423 

mately  follow  the  gas  laws.  On  the  other  hand,  in  1906,  Kah- 
lenberg  found  that  vulcanized  caoutchouc,  so-called  sheet  rubber, 
acts  as  a  semipermeable  membrane  when  used  to  separate  cane 
sugar  solutions  in  pyridine  from  pure  pyridine.  Moreover,  the 
osmotic  pressures  measured  in  pyridine  solutions  thus  far  indi- 
cate that  here  the  gas  laws  do  not  hold  at  all.  Experimental 
osmotic  investigations  are  being  pursued  with  vigor  at  the  pres- 
ent time,  and  it  is  to  be  hoped  that  they  will  cast  more  light 
upon  the  laws  that  regulate  osmotic  pressure  in  the  various 
cases.  The  subject  is  by  no  means  as  simple  as  it  was 
formerly  regarded  by  physical  chemists,  for  specific  selective 
action  on  the  part  of  the  so-called  semipermeable  membrane 
comes  strongly  into  play. 

Solutions  having  the  same  freezing  or  boiling  point  are  fre- 
quently termed  isosmotic  or  isotonic,  for  they  are  said  to  have 
the  same  osmotic  pressure.  The  latter  is  sometimes  computed 
on  the  basis  of  the  assumption  that  the  gas  laws  hold  univer- 
sally for  osmotic  pressures  of  all  solutions,  provided  that  the 
membrane  is  semipermeable.  This  assumption  is,  however,  not 
justified  by  known  experimental  facts.  Indeed,  in  speaking  of 
the  osmotic  pressure  of  a  solution  it  is  always  necessary  to 
specify  what  membrane  separates  that  solution  from  the  pure 
solvent.  A  more  detailed  consideration  of  osmotic  processes 
belongs  to  the  subject  of  physical  chemistry. 

Electrolysis.  —  There  are  two  kinds  of  conductors  of  electricity: 
(1)  those  that  show  no  chemical  change  as  electricity  passes 
through  them,  and  (2)  those  that  conduct  with  concomitant  chem- 
ical decomposition.  To  the  conductors  of  the  first  class  belong  all 
metals,  alloys,  and  graphite.  A  few  other  solids,  like  lead  per- 
oxide, manganese  peroxide,  pyrite,  and  other  native  sulphides 
and  arsenides  also  conduct  slightly  and  probably  without  de- 
composition. Those  conductors  that  are  decomposed  chemi- 
cally by  the  electric  current  are  said  to  conduct  electrolytically. 
They  are  called  electrolytes  or  conductors  of  the  second  class. 
To  this  class  belong  many  metallic  oxides,  hydroxides,  sul- 
phides, chlorides,  nitrates,  sulphates,  carbonates,  and  silicates, 
as  well  as  other  salts  of  metallic  bases,  when  these  compounds 
are  in  molten  condition.  Some  of  these  substances  even  con- 
duct electrolytically  to  an  appreciable  degree  before  they  are 
actually  molten.  So,  for  instance,  at  room  temperatures  a 


424  OUTLINES  OF   CHEMISTRY 

block  of  rock  salt  is  a  non-conductor ;  but  on  heating  it 
strongly,  it  begins  to  conduct  quite  noticeably  before  it  is 
actually  melted.  Again,  the  oxides  of  the  earths  are  non- 
conductors at  room  temperatures,  whereas  at  higher  tempera- 
tures they  conduct,  as  is  well  known  in  the  case  of  the  glower 
of  the  Nernst  lamp.  Water,  alcohols,  ethers,  acids,  esters, 
aldehydes,  ketones,  mercaptans,  sulphur  ethers,  hydrocarbons 
and  their  halogen  substitution  products,  fats,  oils,  waxes,  pitch, 
resin,  caoutchouc,  and  the  chlorides,  bromides,  iodides,  sul- 
phides, oxides,  and  hydrides  of  non-metallic  elements,  when 
pure  are  practically  all  either  exceedingly  poor  conductors  or 
insulators,  whether  they  are  liquefied  or  not.  On  the  other 
hand,  solutions  of  many  acids,  bases,  and  salts  in  water,  liquid 
ammonia,  amines,  liquid  hydrocyanic  acid,  alcohols,  esters, 
ketones,  sulphur  dioxide,  etc.,  are  very  good  electrolytes. 
This  is  particularly  the  case  with  the  ordinary  acids,  salts,  and 
bases  in  aqueous  solutions;  but  it  would  be  quite  wrong  to  think 
that  electrolytic  conduction  is  confined  to  aqueous  solutions,  for 
some  non-aqueous  solutions  conduct  quite  as  well  and  even  better 
than  those  in  which  water  is  the  solvent.  Furthermore,  many 
salts  when  in  the  liquid  condition  or  in  solution  are  insulators. 
So  the  oleates,  palmitates,  and  stearates  of  the  heavy  metals  are 
insulators  whether  solid,  molten,  or  dissolved,  though,  to  be 
sure,  they  do  not  happen  to  dissolve  in  water.  Pure  tetra- 
chloride  of  tin  SnCl4  is  an  insulator  at  all  temperatures  from 
its  freezing  point  to  its  critical  temperature.  The  analogous 
chlorides,  SiCl4,  TiCl4,  and  CC14,  are  also  insulators.  Since  the 
days  of  Michael  Faraday  (1791-1867),  who  first  investigated 
the  subject,  a  large  amount  of  data  on  electrolytic  conductors 
has  been  gathered,  and  yet  no  way  has  been  found  to  foretell  with 
certainty  whether  a  new  compound  will  prove  to  be  a  conductor 
or  not.  The  only  way  is  to  actually  test  the  compound  with  the 
electric  current  itself. 

All  highly  rarefied  gases  conduct  electricity ;  and  in  some 
gases,  like  hydrochloric  acid,  for  instance,  this  conductivity  has 
been  found  to  be  accompanied  with  chemical  decomposition. 
Electrolysis  was  discovered  in  1800  by  Nicholson  and  Carlisle, 
who  decomposed  water  by  means  of  the  electric  current.  This 
new  means  of  effecting  chemical  decomposition  was  studied  by 
Berzelius  and  Hisinger,  Sir  Humphry  Davy,  and  especially  by 


•  SOLUTIONS   AND  ELECTRO-CHEMICAL  THEORIES          425 

Faraday.  The  latter  called  substances  that  can  be  decomposed 
by  the  electric  current  electrolytes;  while  the  process  of  such 
decomposition  he  named  electrolysis.  In  electrolysis  two  plates, 
usually  consisting  of  metal  or  conducting  carbon,  dip  into  the 
electrolyte.  These  plates  Faraday  termed  the  electrodes.  The 
plate  from  which  the  positive  current  passes  into  the  electrolyte 
he  called  the  anode,  and  the  other  electrode,  by  which  the  cur- 
rent leaves  the  electrolyte,  he  called  the  cathode.  It  is  a  re- 
markable fact  that  the  products  of  decomposition  first  appear 
right  on  the  surface  of  the  electrodes.  As  these  products  are 
eliminated  or  produced  at  the  electrodes,  the  corresponding 
material  in  the  electrolyte  is  continually  moving  towards  the 
electrodes.  The  particles  which  thus  move  towards  the  elec- 
trodes during  electrolysis  Faraday  termed  the  ions.  Those 
particles  that  move  towards  the  anode  he  termed  anions,*  and 
those  that  move  towards  the  cathode  the  cations.  To  Faraday, 
the  simplest  conceivable  case  of  electrolysis  consisted  of  two 
silver  electrodes  dipping  into  molten  silver  chloride.  As  a 
current  passes  through  this  electrolyte,  silver  is  deposited  on  the 
cathode,  and  chlorine  is  simultaneously  produced  at  the  anode, 
which  is  thus  at  once  attacked,  forming  silver  chloride.  Thus, 
while  the  cathode  is  increasing  in  thickness,  the  anode  is  wear- 
ing away.  This  process  is,  of  course,  accompanied  by  a  con- 
tinuous movement  of  the  cations,  silver,  toward  the  cathode, 
and  of  the  anions,  chlorine,  toward  the  anode.  If  the  anode 
consisted  of  carbon,  the  chlorine  would  appear  as  a  gas.  Now, 
why  should  molten  silver  chloride  thus  conduct  electrolytically 
and  molten  tetrachloride  of  tin  not?  This  is  a  question  which 
we  cannot  yet  answer,  any  more  than  we  can  tell  why  a  piece 
of  silver  conducts  electricity  and  a  piece  of  sulphur  is  an  in- 
sulator. In  solutions,  electrolysis  is  precisely  the  same  as  in 
molten  electrolytes,  though,  to  be  sure,  the  products  eliminated 
at  the  electrodes  in  solutions  react  with  the  electrolyte  more 
frequently  because  of  the  more  complex  nature  of  the  latter. 
So,  on  electrolysis  of  an  aqueous  sodium  sulphate  solution, 
hydrogen  appears  at  the  cathode  and  oxygen  at  the  anode. 
This  is  because  the  sodium  eliminated  at  the  cathode  at  once 
reacts  with  the  water,  forming  caustic  soda  and  hydrogen,  which 
is  consequently  a  secondary  product.  If  a  mercury  cathode  is 
used,  the  sodium  liberated  dissolves  in  the  latter  and  no  hydro- 


426  OUTLINES   OF   CHEMISTRY 

gen  appears.  At  the  anode  SO4  is  liberated,  which,  however, 
at  once  reacts  with  water,  yielding  sulphuric  acid  and  oxygen, 
which  is  consequently  of  secondary  origin,  like  the  hydrogen 
at  the  cathode. 

In  1833  Faraday  demonstrated  that  the  passage  of  the  electric 
current  through  an  electrolyte  is  always  accompanied  by  the  ap- 
pearance of  decomposition  products  at  the  electrodes,  and  that 
the  amount  of  such  decomposition  is  proportional  to  the  current. 
Furthermore,  he  found  that  chemically  equivalent  amounts  of 
substances  are  separated  out  from  different  electrolytes  by  the 
same  amount  of  current.  These  facts  are  commonly  known  as 
Faraday's  law.  They  are  fundamental  in  all  electrolytic  work. 
Faraday  tested  his  law  on  molten  electrolytes  as  well  as  on 
aqueous  solutions,  and  its  validity  has  since  been  confirmed 
by  many  careful  investigations.  In  1900  Kahlenberg  showed 
that  the  law  holds  also  in  non-aqueous  solutions.  A  current 
of  1  ampere  will  effect  the  deposition  of  a  gram-equivalent 
of  a  substance  in  96,540  seconds ;  that  is  to  say,  whenever 
96,540  coulombs  of  electricity  pass  through  an  electrolyte,  a 
gram-equivalent  of  decomposition  product  is  deposited  at  each 
electrode.  The  quantity  96,540  coulombs  is  known  as  the  con- 
stant of  Faraday's  law ;  it  is  now  generally  called  a  faraday. 

Electrolytic  Theories.  The  Grotthus  Theory.  —  The  first  ex- 
planation of  the  process  of  electrolysis  was  made  by  Grotthus, 
in  1805.  He  assumed  that  each  molecule  of  an  electrolyte 
possesses  electrical  polarity  similar  to  that  of  a  magnet,  and 
that  these  molecules  are  irregularly  distributed  throughout 
the  electrolyte.  On  closing  the  electric  circuit,  the  +  poles 
of  the  molecules  would  be  attracted  by  the  —  electrode,  and  the 
—  poles  of  the  molecules  by  the  +  electrode.  Thus  the  mole- 
cules in  the  electrolyte  would  all  arrange  themselves  with  their 
+  poles  directed  toward  the  —  electrode,  and  their  —  poles 
toward  the  +  electrode.  The  molecules  in  actual  contact  with 
the  electrodes  were  then  conceived  as  decomposed  by  the  attrac- 
tion of  the  -|-  electrode  for  the  —  part  of  the  molecule,  and  of 
the  —  electrode  for  the  +  part  of  the  molecule.  So  with  molten 
silver  chloride  as  electrolyte  between  two  platinum  electrodes 
(Fig.  145),  after  closing  the  circuit,  the  arrangement  would  be 
as  shown  in  line  (1).  A  moment  later  silver  would  be  depos- 
ited at  Q  and  chlorine  at  A,  and  line  (2)  would  represent  the 


SOLUTIONS  AND  ELECTRO-CHEMICAL  THEORIES 


427 


condition  of  the  remaining  molecules.  A  general  decomposition 
all  along  the  line  would  then  take  place,  so  that  the  arrange- 
ment in  line  (3)  would  result.  Then  these  molecules  would 
turn  on  an  axis,  their  —  poles  heading  toward  the  +  electrode, 


FIG.  145. 

and  the  +  poles  toward  the  —  electrode,  as  in  line  (1),  after 
which  the  whole  process  would  repeat  itself. 

Faraday's  View.  —  This  mechanical  explanation  of  Grotthus 
held  its  own  for  over  half  a  century,  though  Faraday  protested 
that  any  explanation  that  the  molecules  were  rent  asunder  by 
an  attraction  of  the  electrodes  for  the  molecules  was  untenable. 
According  to  Faraday,  all  that  can  be  maintained  is  that  the 
electric  current  acts  as  an  axis  of  force,  ejecting  the  decomposi- 
tion products  at  the  electrodes,  which  simply  serve  as  doors  for 
the  exit  of  the  products. 

Clausius's  Theory.  —  In  1856  Robert  Clausius  called  special 
attention  to  the  fact  that  the  theory  of  Grotthus  does  not 
explain  why  a  small  electro-motive  force  suffices  to  send  an 
appreciable  current  through  an  electrolyte,  though,  to  be  sure, 
such  current  might  pass  for  but  a  short  time.  It  was  at  the 
time  when  the  kinetic  theory  of  gases  and  the  mechanical 
theory  of  heat  were  taking  form.  In  these  theories  Clausius 
was  interested.  He  consequently  naturally  assumed  the  mole- 
cules of  an  electrolyte  to  be  in  motion  due  to  heat,  just  as  the 
molecules  of  a  gas  are  supposed  to  be  in  motion.  He  further 
stated  that  sometimes  some  of  the  molecules  would  collide  in 
such  a  way  that  the  basic  and  acid  radicals  of  two  different 
molecules  would  unite  to  form  a  new  molecule,  thus  leaving  a 
basic  and  an  acid  radical  in  an  uncombined  or  free  state  for 


428  OUTLINES  OF   CHEMISTRY 

a  moment.  At  any  instant,  then,  there  would  be  a  certain 
number  of  free  positive  basic  radicals  and  free  negative  radi- 
cals, and  these  would  be  separated  at  the  electrodes  by  electrical 
attraction,  as  Grotthus  explained.  However,  Clausius  did  not 
assume  that  the  molecules  ever  arranged  themselves  definitely 
in  the  electrolyte.  He  thought  that  the  electric  current  merely 
directed  the  general  trend  of  the  decomposition  within  the 
electrolyte.  Moreover,  Clausius  held  that  the  free,  or  unde- 
composed,  parts  of  an  electrolyte  would  at  any  moment  not 
amount  to  more  than  a  small  fraction  of  1  per  cent.  This 
distinguishes  Clausius's  theory  from  that  of  Arrlienius. 

Arrhenius's  Theory.  —  The  theory  of  Arrhenius,  also  known 
as  the  theory  of  electrolytic  dissociation,  and  frequently  nowa- 
days designated  as  the  ionic  theory,  is  founded  upon  a  supposed 
connection  between  the  vapor  tensions,  boiling  points,  or  freez- 
ing points  of  dilute  solutions  of  electrolytes,  on  the  one  hand, 
and  their  electrolytic  conductivity  on  the  other  hand.  In  1887 
Arrhenius  published  a  series  of  data  intended  to  show  that 
dilute  solutions  of  electrolytes  lower  the  freezing  point  to  a 
much  greater  extent  than  similar  solutions  of  non-electrolytes 
containing  equimolecular  quantities,  and  that  consequently  the 
solutions  that  conduct  the  current  must  contain  relatively  more 
dissolved  molecules  ;  that  is,  the  latter  must  be  dissociated. 
Because  this  dissociation  is  supposed  to  occur  in  electrolytes, 
it  has  been  named  electrolytic  dissociation.  Arrhenius  assumes 
that  when  an  acid,  base,  or  salt  is  dissolved,  yielding  a  solution 
that  conducts  the  current,  the  molecules  of  the  dissolved  sub- 
stance are  by  the  very  act  of  solution  decomposed  into  part 
molecules,  which  are  charged  with  electricity,  thus  :  — 


NaOH±N+a+O~H. 


H2SO4  ;£  H  +  HSO4  or  H  +  H  +  SO4. 


Ag  +  NO3,  etc. 

In  each  case  the  dissociation  from  left  to  right  is  supposed  to 
be  complete  in  infinitely  dilute  solutions.     At  finy  finite  concen- 


SOLUTIONS   AND   ELECTRO-CHEMICAL  THEORIES 


429 


tration  the  solution  contains  a  certain  percentage  of  undissociated 
molecules  and  a  certain  percentage  of  charged  part  molecules. 
The  latter  according  to  Arrhenius  are  the  only  particles  in  the 
solution  that  are  active  in  conducting  the  current.  He  calls 
these  part  molecules  the  ions,  and  conceives  each  gram-equiva- 
lent of  cations  charged  with  96,540  coulombs  of  positive  electric- 
ity and  each  gram-equivalent  of  anions  charged  with  the  same 
amount  of  negative  electricity,  postulating  further  that  any 
conducting  solution  contains  equivalent  amounts  of  cations  and 
anions.  It  will  be  observed  that  in  the  Arrhenius  theory  the 
word  ion  is  used  in  a  different  sense  from  that  proposed*  by 
Faraday,  who  regarded  the  ions  as  the  substances  that  migrate 
toward  the  electrodes  during  actual  electrolysis,  not  as  part 
molecules  charged  with  electricity  which  are  at  all  times  present 
in  an  electrolyte.  The  explanation  of  the  passage  of  the  elec- 
tric current  through  an  electrolyte  on  the  basis  of  Arrhenius's 
theory  is  simple,  consisting  merely  of  the  movement  of  the  free 
charged  cations  and  anions  toward  the  cathode  and  anode 
respectively,  under  the  influence  of  the  impressed  electro-motive 
force  (Fig.  146).  This  explanation  is  then  essentially  the  same 


IB 


E> 


(9 
0GGJGG   <QGG 

6 


S> 


Fi<a.  146. 

as  that  given  by  Clausius.  Indeed,  the  main  difference  between 
the  Clausius  and  Arrhenius  theories  is  that  the  latter  assumes 
the  presence  of  a  very  much  larger  percentage  of  dissociation, 
and  claims  this  may  be  computed  from  either  conductivity  or 
freezing  point  or  boiling  point  measurements  of  dilute  solutions. 
According  to  Arrhenius's  theory  all  the  physical,  chemical, 
and  physiological  properties  of  solutions  that  are  electrolytes  are 


430  OUTLINES   OF  CHEMISTRY 

determined  by  the  properties  of  the  undissociated  molecules  of 
solute  and  those  of  the  ions.  The  latter  are  the  chief  factor  in 
dilute  solutions  where  dissociation  has  often  progressed  to  the 
extent  of  80  per  cent  or  more.  Thus,  the  adherents  of  this 
theory  hold  that  copper  ions  are  blue,  cobalt  ions  are  red, 
MnO4  ions  are  purple,  sodium  ions  are  colorless,  etc.,  since 
aqueous  solutions  of  copper  salts  are  blue,  those  of  cobalt  are 
red,  those  of  permanganates  purple,  and  those  of  sodium  chlo- 
ride, etc.,  colorless.  Again,  the  sour  taste  and  other  acid 
properties  of  electrolytic  solutions  of  acids  would  be  due  to 
hydrogen  ions,  and  the  alkalinity  of  caustic  alkali  solutions  to 
hydroxyl  ions.  Indeed,  in  terms  of  the  theory  of  electrolytic  disso- 
ciation an  acid  would  be  defined  as  a  substance  capable  of  yielding 
hydrogen  ions,  while  a  base  would  be  a  substance  capable  of  yield- 
ing hydroxyl  ions.  Moreover,  the  act  of  neutralization  in  dilute 
solutions  would  consist  essentially  of  the  union  of  hydrogen  and 
hydroxyl  ions  to  form  undissociated  water,  thus:  — 

K,  OH  +  H,  Cl  ==  K,  01+  H2O. 

Since  K  and  Cl  appear  on  both  sides  of  the  equation,  the  latter 
might  even  be  written:  — 

0~H  +  H  =  H20. 

In  the  chapter  on  thermochemistry  it  has  been  stated  that  the 
heat  of  neutralization  of  strong  acids  by  strong  bases  is  approxi- 
mately the  same  for  all.  This  fact  clearly  would  readily  be 

explained  by  the  above  assumption  that  the  neutralization  in  all 

+ 
cases  consists  essentially  of  a  union  of   H  and  OH.     On  the 

other  hand,  it  is  to  be  noted  that  the  heats  of  neutralization  of 
some  of  the  weaker  acids,  like  oxalic  and  acetic  acids,  are  nearly 
the  same  as  that  of  hydrochloric  acid  ;  though  according  to  the 
theory,  the  latter  is  far  more  highly  dissociated. 

Again,  the  fact  that  two  solutions  of  neutral  salts  which  form 
no  precipitate  when  mixed  also  exhibit  no  change  of  temperature 
on  being  poured  together  (so-called  law  of  thermoneutrality  of 
Hess)  is  readily  explained  by  Arrhenius's  hypothesis,  for  there 
would  be  no  action  to  occasion  a  thermal  effect.  However, 
when  similar  salt  solutions  that  are  non-electrolytes  are  mixed, 
there  is  also  no  thermal  effect. 

Precipitation  in  electrolytic  solutions  by  double  decomposi- 


SOLUTIONS  AND   ELECTRO-CHEMICAL  THEORIES          431 

tion  is  explained  on  the  basis  of  the  Arrhenius  theory  by  saying 
that  certain  ions  meet  to  form  insoluble  compounds.  The  pre- 
cipitation of  silver  chloride  when  silver  nitrate  and  common  salt 
solutions  are  brought  together  would  be  expressed  as  follows  :  — 


Ag,  NO3  +  Na,  Cl  =  Na,  NO3  +  AgCl. 

In  other  words,  the  chlorine  ions  and  silver  ions  meet  to  form 
neutral  silver  chloride,  which  is  difficultly  soluble.  Since  all 
conducting  chloride  solutions  are  supposed  to  contain  free 
chlorine  ions,  it  is  clear  that  these  would  all  precipitate  silver 
chloride  from  solutions  of  silver  salts  that  contain  silver  ions. 
Similarly,  the  fact  that  barium  sulphate  is  precipitated  when  a 
solution  of  a  barium  salt  and  one  of  a  soluble  sulphate  are 
brought  together  finds  a  ready  explanation,  etc.  At  one  time 
the  view  was  even  advocated  that  all  chemical  reactions,  or  at 
least  those  that  take  place  instantaneously,  occur  only  between 
free  ions,  that  is  to  say,  they  occur  only  in  electrolytes. 

Now  the  fact  is  that  all  chemical  reactions  known  to  occur  in 
solutions  that  are  electrolytes  can  be  reproduced,  as  to  type,  in 
solutions  that  are  the  best  of  insulators.  So,  for  instance,  from 
copper  oleate  solutions  in  hydrocarbon  oils,  brown  cupric  chlo- 
ride is  instantly  precipitated  by  adding  any  of  the  following 
chlorides,  also  dissolved  in  the  same  hydrocarbon  :  HC1,  PC13, 
SnCl4,  SiCl4,  SbCl3,  etc.,  though  these  solutions  are  all  non- 
electrolytes.  Lead  will  precipitate  copper  from  insulating 
copper  oleate  solutions  just  as  zinc  precipitates  copper  from 
aqueous  copper  sulphate  solutions.  Furthermore,  the  colors 
exhibited,  for  example,  by  solutions  of  copper  oleate,  cobalt 
oleate,  nickel  oleate,  etc.,  each  dissolved  in  a  hydrocarbon 
like  benzene,  toluene,  etc.,  are  entirely  similar  to  the  colors 
of  aqueous  electrolytic  solutions  of  copper  and  cobalt  salts 
respectively.  In  fact  all  the  physical  and  chemical  properties 
exhibited  by  salt  solutions  that  are  electrolytes  can  be  duplicated 
in  salt  solutions  that  are  insulators,  except  the  phenomena  of  elec- 
trolysis themselves. 

But  the  distinguishing  feature  of  Arrhenius's  theory  is  the 
claim  that  there  is  a  quantitative  relation  between  lowering  of 
the  freezing  point  or  elevation  of  the  boiling  point  and  the  elec- 
trical conductivity  of  solutions.  It  must  be  stated  that  the 
numerous  cases  thus  far  adduced  to  support  this  claim  are  not 


432  OUTLINES   OF   CHEMISTRY 

at  all  conclusive,  for  they  show  variations  that  are  far  beyond 
the  limits  of  experimental  errors.  Furthermore,  solutions  of 
magnesium  sulphate  as  well  as  those  of  all  the  other  vitriols 
are  good  electrolytes,  whereas  according  to  the  freezing  points 
of  their  solutions  they  ought  to  be  non-electrolytes.  Again, 
soap  solutions  boil  at  practically  the  same  temperature  as  water, 
yet  they  conduct  electricity  well.  And  so  numerous  other 
cases  might  be  cited,  the  field  of  non-aqueous  solutions  being 
especially  replete  with  such.  It  should  be  stated  that  the 
behavior  of  electrolytes  is  in  general  not  in  harmony  with  the 
law  of  mass  action,  as  ought  to  be  the  case  if  Arrhenius's 
theory  were  tenable. 

A  more  detailed  account  of  this  interesting  theory  cannot  be 
given  here.  Suffice  it  to  say  that,  taking  all  known  facts  into 
consideration,  the  theory  of  Arrhenius  appears  untenable.  This 
view  was  also  clearly  voiced  by  Mendeleeff  in  the  last  edition 
of  his  great  work  on  the  "  Principles  of  Chemistry,"  when, 
referring  to  the  theory  of  electrolytic  dissociation,  he  said,  "I 
do  not  consider  the  hypothesis  in  question  to  be  in  accordance 
with  fact,  and  therefore  refrain  from  giving  a  detailed  exposition 
of  it  in  this  work."  The  reader  will  have  no  difficulty  in  com- 
prehending books  that  still  use  the  nomenclature  of  the  theory 
of  electrolytic  dissociation  by  remembering  that  the  term  ion  as 
used  in  expressing  chemical  changes  means  the  same  as  atom  or 
radical. 

The  electron  theory  considers  electricity  itself  to  be  material 
in  character  and  to  consist  of  corpuscles  or  electrons  that  weigh 
about  0.0005  as  much  as  a  hydrogen  atom.  This  theory  has 
developed  from  a  study  of  radium  rays  and  the  discharge  of 
electricity  through  rarefied  gases.  The  electrons  are  considered 
to  be  negative  electricity  itself.  Positive  electrons  appear  to 
be  much  more  difficult  to  isolate.  J.  J.  Thomson  has  recently 
attempted  to  construct  a  theory  that  the  atoms  of  the  various 
elements  are  composed  entirely  of  electrons,  and  has  shown  that, 
on  the  basis  of  such  an  assumption,  the  properties  of  the  ele- 
ments would  exhibit  periodicity  as  indicated  by  the  periodic 
system.  The  electron  theory  has  not  yet  been  tested  as  to  its 
value  in  the  study  of  chemical  changes. 

Electric  Batteries  are  contrivances  for  converting  chemical 
energy  into  electrical  energy.  So  when  a  zinc  plate  and  a 


SOLUTIONS  AND   ELECTRO-CHEMICAL  THEORIES 


433 


graphite  plate  are  dipped  into  dilute  sulphuric  acid,  as  shown  in 
Fig.  147,  zinc  dissolves  and  hydrogen  is  liberated  on  the  graph- 
ite ;  while  at  the  same  time  an  electric  current  passes  through 
the  solution  from  the  zinc  to  the  graphite 
and  from  the  latter  through  the  wire  to 
the  zinc.     Some  hydrogen  also  appears  at 
the  zinc,  though  in  very  small  amounts  if 
the  zinc  is  pure  or  if  it  is  amalgamated. 
The  chemical  action  that  takes  place  is: — 

Zn  +  H2SO4  =  ZnSO4  +  H2. 

But  the  zinc  dissolves  at  one  plate  and  the 
hydrogen  appears  at  the  other,  while  in 
the  middle  of  the  electrolyte  there  is  no 
visible  change  for  some  time.  The  ex- 
planation of  the  passage  of  the  current  FlG  147 
through  the  electrolyte  is  the  same  as  in 

the  case  of  electrolysis.  In  a  battery,  however,  no  external 
electro-motive  force  is  used  to  force  the  current  through  the 
electrolyte.  Batteries  develop  electro-motive  force  of  their 
own  because  of  the  chemical  affinity  that 
comes  into  play ;  and  the  electrical  energy 
developed  by  batteries  comes  from  the  energy 
of  the  chemical  changes  that  take  place  while 
the  battery  is  in  action.  The  cell  above 
mentioned  will  exhibit  an  E.  M.  F.  of  about 
1.3  volts,  which  on  closed  circuit  soon  drops 
rapidly  because  of  the  accumulation  of  hy- 
drogen on  the  graphite,  for  this  hydrogen- 
laden  plate  produces  a  counter  E.  M.  F. 
Any  two  different  conductors  of  the  first 
|  :  class  dipping  into  an  electrolyte  will  show  an 
I1  E.  M.  F.  when  connected  with  a  voltmeter  as 

shown  in  Fig.  148.  A  complete  considera- 
tion of  electric  batteries  belongs  to  the  sub- 
ject of  electro-chemistry  and  cannot  be 
entered  into  here. 

The  ordinary  batteries  used  for  ringing 
doorbells  consist  of  zinc  and  carbon  dipping  into  a  concentrated 
solution  of  ammonium  chloride.  As  the  battery  acts,  zinc  is 

2F 


FIG.  148. 


434 


OUTLINES    OF  CHEMISTRY 


dissolved,  forming  zinc  chloride,  while  hydrogen  is  liberated  on 
the  carbon,  thus  :  — 

Zn  +  2  NH4C1  +  2  H2O  =  ZnCl2  +  2  NH4OH  +  H2. 
The  ammonium  hydroxide  forms  at  the  carbon.  As  these  bat- 
teries are  used  only  occasionally,  and  then  only  for  a  short  time, 
this  hydrogen  ordinarily  escapes  while  the  battery  is  at  rest,  thus 
prolonging  its  life  and  usefulness.  The  dry  batteries  in  use  com- 
monly consist  of  zinc,  ammonium  chloride  solution,  and  carbon, 
the  latter  being  surrounded  with  coarsely  powdered  manganese 
dioxide  which  serves  to  oxidize  the  hydrogen  liberated.  These 
batteries  are  not  perfectly  dry,  as  their  name  would  indicate,  but 
they  contain  enough  plaster  of  Paris  to  solidify  their  contents. 
Another  battery,  which  frequently  is  used  in  telegraphic 
work,  consists  of  zinc  dipping  into  dilute  sulphuric  acid,  and  cop- 
per surrounded  by  a  saturated 
solution  of  copper  sulphate. 
The  two  solutions,  being  of  dif- 
ferent density,  are  kept  separate 
by  gravity  (Fig.  149).  The 
battery  must  be  kept  in  use  all 
the  time,  however,  or  the  solu- 
tions will  diffuse  into  each  other, 
and  the  copper  sulphate  solution 
will  reach  the  zinc  and  react 
with  it,  forming  zinc  sulphate 
and  copper,  which  by  coating 
the  zinc  would  spoil  the  battery. 
The  E.  M.  F.  of  this  battery, 
known  as  the  blue  cup  battery, 
is  about  1.1  volts. 

The  ordinary  storage  battery 
consists  of  two  lead  plates,  one 
of  which  is  coated  with  lead 

peroxide,  dipping  into  a  solution  of  sulphuric  acid  of  1.2  specific 
gravity.  As  the  battery  acts,  lead  sulphate  is  formed.  The 
complete  change  may  be  represented  thus :  — 

Pb  +  H2S04  =  PbS04  +  2  H,  and 
PbO2  +  2  H  +  H2S04  =  PbS04  +  2  H20  ;  or 
Pb  +  2  H2S04  +  Pb02  =  2  PbS04  +  2  H20. 


FIG.  149. 


SOLUTIONS  AND   ELECTRO-CHEMICAL   THEORIES          435 

At  the  cathode  lead  dissolves,  while  at  the  anode  the  hydrogen 
liberated  reduces  the  lead  peroxide,  which  is  then  acted  upon 
by  sulphuric  acid,  forming  lead  sulphate.  The  E.  M.  F.  of  the 
storage  cell  is  about  2  volts.  When  in  use  its  voltage  decreases. 
When  the  E.  M.  F.  has  run  down  to  1.8  volts,  the  battery 
should  be  recharged.  This  is  done  by  passing  a  current  from 
a  dynamo  through  the  cell  in  the  opposite  direction  from  that 
produced  by  the  battery  when  in  action.  In  the  charging  pro- 
cess the  reactions  just  given  are  reversed ;  that  is,  lead  is  depos- 
ited at  the  cathode  and  lead  peroxide  is  formed  at  the  anode. 

Electro-chemical  Series  of  the  Metals.  —  If  in  the  battery  zinc, 
dilute  sulphuric  acid,  carbon  (Fig.  147),  the  zinc  be  replaced 
successively  by  iron,  copper,  silver,  and  gold,  the  voltage  will 
diminish.  The  order  in  which  the  metals  are  mentioned  is  the 
same  as  that  in  which  they  will  replace  one  another  in  solutions 
of  their  salts.  All  of  the  metals  may  thus  be  arranged  in  a 
series,  beginning  with  the  most  basic,  and  ending  with  the  least 
basic ;  or,  as  it  is  sometimes  stated,  beginning  with  the  most 
electro-positive  metal  and  ending  with  the  least  electro-positive. 
Such  a  series  is  called  the  electro-chemical  series  of  the  metals. 
The  series  varies  somewhat  for  different  solutions,  but  the  usual 
order  for  the  common  metals  is  about  as  follows  :  K,  Na,  Ba,  Sr, 
Ca,  Mg,  Al,  Mn,  Zn,  Cd,  Fe,  Co,  Ni,  Pb,  Bi,  Sb,  Sn,  Cu,  Hg,  Ag, 
Pt,  Au.  This  is  also  called  arranging  the  metals  in  the  order 
of  their  electrolytic  solution  tensions  by  those  who  think  of  the 
metals  as  having  a  tension  or  tendency  to  form  ions  in  the  sense 
of  the  theory  of  Arrhenius. 

The  ease  with  which  the  metals  are  separated  from  electro- 
lytes by  electrolysis  is  in  the  reverse  order  of  that  given  above. 


CHAPTER  XXV 

COPPER,    SILVER,    AND   GOLD 

COPPER,  silver,  and  gold  occur  in  nature  in  the  uncombined 
state,  and  consequently  have  been  known  to  man  since  earliest 
times.  These  metals  have  a  high  specific  gravity,  are  very 
malleable  and  ductile,  and  most  excellent  conductors  of  heat 
and  electricity.  The  atomic  weights  of  the  elements  of  this 
group  are  relatively  high:  Cu,  63.57;  Ag,  107.88;  Au,  197.2. 
Chemically  these  metals  are  rather  inert,  and  their  chemical 
activity  decreases  as  the  atomic  weight  increases.  Like  mer- 
cury, copper  forms  two  series  of  compounds,  the  cuprous,  in 
which  copper  is  univalent,  and  the  cupric,  in  which  it  is  biva- 
lent. Silver  is  practically  always  univalent  in  its  compounds. 
Gold,  on  the  other  hand,  forms  aurous  compounds,  in  which 
the  metal  is  univalent,  and  auric  compounds,  in  which  it  is 
trivalent.  The  compounds  of  copper,  silver,  and  gold  in  which 
these  metals  are  univalent  are  analogous  to  those  of  the  alkali 
metals,  which  are  also  univalent.  On  the  other  hand,  copper 
in  cupric  compounds,  in  which  it  is  bivalent,  is  analogous  to 
mercury  in  mercuric  compounds.  Copper  and  gold  are  the 
only  colored  metals  known.  Copper,  silver,  and  gold  are  used 
in  all  civilized  countries  in  making  coins. 

Occurrence,  Metallurgy,  and  Properties  of  Copper.  —  In  the 
uncombined  condition  'copper  is  found  in  large  quantities  near 
Lake  Superior.  It  also  occurs  in  the  Urals,  in  Sweden, 
Japan,  and  China.  Large  amounts  of  chalcocite  or  copper 
glance  Cu2S,  chalcopyrite  or  copper  pyrites  Cu2S  •  Fe2S3,  and 
bornite  (Cu2S)3Fe2S3  occur  in  Montana,  where  they  are  mined 
and  smelted  for  copper.  Other  important  copper  ores  are 
ruby  copper  Cu2O,  malachite  CuCO3  •  Cu(OH)2,  and  azurite 
Cu(OH)2  •  2  CuCO3.  Malachite  occurs  especially  in  Siberia. 
In  extremely  small  amounts  copper  compounds  are  also  some- 
times found  in  plants  and  animals.  Thus,  in  plants  growing 
in  copper- bearing  regions,  copper  is  frequently  met;  similarly 

436 


COPPER,    SILVER,   AND   GOLD  437 

in  oysters ;  also  in  the  feathers  of  some  birds,  like  those  of  the 
genus  turacus. 

In  extracting  copper  from  its  ores,  the  process  is  simple 
when  the  ores  are  oxides  or  carbonates,  for  then  all  that  is  nec- 
essary is  to  heat  the  ore  with  coke  or  coal  in  a  blast  furnace ; 
the  reaction  involved  is  :  — 

Cu2O  4-  C  =  2Cu  +  CO. 

But  if  the  ores  contain  sulphides,  which  is  commonly  the  case, 
the  process  is  much  more  difficult,  for  iron  and  other  impurities, 
like  lead,  arsenic,  and  antimony,  besides  sulphur,  must  be  elim- 
inated. For  ores  rich  in  copper  a  dry  process  is  used,  whereas 
for  ores  that  have  a  low  copper  content  the  wet  process  is  com- 
monly adopted. 

In  the  dry  process,  the  ores  are  roasted  to  convert  most  of 
the  sulphides  into  oxides.  In  this  way,  sulphur  burns  to  sul- 
phur dioxide,  and  iron  to  iron  oxides  to  a  large  extent.  The 
mass  is  then  mixed  with  carbon  and  silicates  rich  in  silica,  and 
heated  in  a  blast  furnace.  Thus  iron  enters  the  slag  as  a  sili- 
cate, which  floats  on  top  and  is  tapped  off.  In  the  process, 
copper  oxides  are  reduced,  the  copper  formed  uniting  with  any 
sulphides  still  present,  forming  a  heavy  liquid  which  accumu- 
lates at  the  bottom  of  the  furnace.  This  liquid  is  run  off,  fre- 
quently into  water  so  as  to  granulate  the  product,  which  is 
called  copper  matte,  and  consists  mainly  of  copper  and  iron 
sulphides  containing  about  33  per  cent  copper.  This  copper 
matte  is  then  roasted,  and  again  melted  with  silicates  and  carbon. 
Thus  more  iron  is  removed  as  silicate,  and  copper  and  other  ox- 
ides are  reduced  to  metal  containing  90  to  95  per  cent  copper. 
This  product  is  then  melted  on  the  hearth  of  a  furnace  in  a 
current  of  air,  whereby  metals  like  iron,  arsenic,  lead,  and 
antimony  are  oxidized  before  copper,  and  these  oxides  either 
pass  off  as  vapors  or  float  on  the  surface  and  are  removed. 
The  copper  is  also  slightly  oxidized,  and  -this  oxide  reacts  with 
any  cuprous  sulphide  left  over,  forming  copper  and  sulphur 
dioxide,  thus :  — 

Cu2S  +  2  CuO  =  4  Cu  +  SO2. 

Finally,  by  adding  carbon,  any  further  copper  oxide  is  reduced 
to  metal.  Bessemer  converters  (which  see)  are  sometimes 
used  in  effecting  the  oxidation  of  the  ores  or  the  copper  matte. 


438  OUTLINES  OF  CHEMISTRY 

Air  containing  sand  or  other  finely  divided  silicates  is  blown 
into  these  converters  containing  the  molten  matte.  Thus  the 
process  above  described  is  conducted  in  practically  one  opera- 
tion. 

In  the  wet  process,  ores  low  in  copper,  like  oxidized  pyrites 
residues  from  sulphuric  acid  factories,  are  roasted  with  common 
salt,  thus  forming  cupric  chloride,  which  is  leached  out  with 
water.  From  the  clear  solution,  copper  is  precipitated  by  scrap 
iron,  thus :  — 

CuCl2  +  Fe  =  FeCl2  +  Cu. 

Or  the  ores  are  moistened  with  water  and  exposed  to  the  air ; 
thus  ferric  sulphate  gradually  forms,  in  a  solution  of  which 
copper  is  soluble.  This  solution  is  then  treated  with  scrap 
iron  to  precipitate  the  copper.  The  supernatant  solution,  con- 
sisting essentially  of  ferrous  sulphate,  is  again  poured  upon 
fresh  ore,  and  so  on. 

Electrolysis  is  used  in  copper  refining.  The  impure  copper 
from  the  above  processes  is  melted  and  cast  into  thick  plates. 
These  serve  as  anodes,  a  copper  sulphate  solution  is  used  as  the 
electrolyte,  and  thin  copper  plates  serve  as  cathodes.  The 
copper  is  easily  eroded  from  the  anode  and  deposited  on  the 
cathode.  As  the  process  virtually  consists  of  simply  trans- 
ferring copper  from  anode  to  cathode,  an  E.  M.  F.  below  one 
volt  is  sufficient  to  keep  the  current  flowing.  The  deposited 
copper  is  stripped  off  from  the  cathodes  when  they  have  become 
sufficiently  thick.  As  the  impure  copper  anodes  are  dissolving, 
the  impurities,  like  copper  sulphide,  silver,  gold,  bismuth,  and 
lead,  fall  to  the  bottom  of  the  vats,  forming  a  thick  mud  or 
slime,  from  which  the  two  precious  metals  mentioned  are  sep- 
arated. Electrolytically  refined  copper  is  in  a  high  state  of 
purity.  Its  malleability,  ductility,  and  electrical  conductance 
are  much  greater  than  that  of  impure  copper. 

In  1910  the  United  States  produced  540,079  tons  of  copper, 
which  is  somewhat  over  half  of  the  world's  output  for  the  year. 

Copper  is  a  rather  hard,  tough,  very  ductile  and  malleable 
metal  of  reddish  brown  color  and  high  metallic  luster.  It  may 
be  crystallized  in  regular  octahedra  or  cubes.  In  dry  air  it  re- 
mains unchanged.  In  moist  air  it  gradually  becomes  coated 
with  a  greenish  basic  copper  carbonate  known  as  verdigris.  The 
specific  gravity  of  copper  is  8.94.  Its  melting  point  is  1084°, 


COPPER,    SILVER,   AND   GOLD  439 

and  its  boiling-point  1500°.  Just  below  its  melting  point 
copper  becomes  so  brittle  that  it  may  be  pulverized.  At  red 
heat  it  may  be  welded.  Molten  copper  absorbs  gases  like  car- 
bon monoxide,  hydrogen,  and  sulphur  dioxide,  which  are  given 
off  with  some  violence  on  cooling.  Copper  is  readily  oxidized 
by  heating  it  in  the  air.  It  imparts  a  green  color  to  the  Bunsen 
flame.  In  dilute  sulphuric  acid,  copper  is  insoluble,  but  the  hot 
concentrated  acid  attacks  it,  forming  copper  sulphate,  sulphur 
dioxide,  and  water.  In  dilute  nitric  acid  copper  is  readily 
soluble :  — 

3  Cu  +  8  HN03  =  3  Cu(N03)2  +  4  H2O  +  2  NO. 

Cold  hydrochloric  acid  is  practically  without  action  on  copper, 
but  the  hot  concentrated  acid  attacks  the  metal  somewhat, 
hydrogen  and  cuprous  chloride  being  formed.  Ammonia  water 
slowly  dissolves  copper  in  contact  with  air.  The  deep  blue 
solution  thus  obtained,  known  as  Schweitzer's  reagent,  is  a  good 
solvent  for  cellulose.  Many  dilute  acids  also  gradually  act  on 
copper,  in  presence  of  oxygen  or  air,  which  fact  is  important, 
since  copper  compounds  are  poisonous  and  copper  is  often  used 
in  making  cooking  pots,  kettles,  and  other  utensils.  Large 
quantities  of  copper  are"  used  for  electric  wires  and  cables,  for 
covering  the  bottoms  of  ships,  for  coins,  for  roofing  and  archi- 
tectural ornamentation,  for  electroplating,  and  for  a  number  of 
important  alloys.  Copper  salts  are  important  as  germicides  in 
agriculture.  From  solutions  of  its  salts,  copper  is  precipitated 
by  magnesium,  zinc,  iron,  phosphorus,  and  various  other  reduc- 
ing agents. 

Alloys  of  Copper.  — Most  of  the  copper  produced  is  used  in  the 
form  of  alloys.  Pure  copper  does  not  make  good  castings,  for 
on  cooling  it  contracts  and  does  not  fill  the  molds.  Many  alloys 
of  copper,  however,  are  excellent  for  making  castings,  and  they 
are  constantly  of  very  great  importance.  The  alloys  are  made 
by  melting  the  metals  together  in  the  proportions  required. 
Copper  readily  forms  alloys  with  most  of  the  metals.  Among 
the  most  important  of  these  alloys  are  the  following: — 

Brass  is  a  golden  yellow  alloy  consisting  of  2  parts  copper  arid 
1  part  zinc.  Other  proportions  may  be  used,  however;  thus 
5  parts  copper  and  1  part  zinc  yield  a  reddish  brass  known  as 
tombac,  Dutch  brass,  or  Dutch  metal.  Muntz  metal  contains 


440  OUTLINES   OF   CHEMISTRY 

3  parts  copper  and  1  part  zinc.  The  larger  the  content  of  zinc, 
the  paler  the  color  of  the  brass  obtained  ;  the  varieties  of  brass 
on  the  market  commonly  contain  from  18  to  50  per  cent  zinc. 
Brass  is  harder  than  copper.  It  is  nevertheless  malleable  and 
ductile  and  makes  good  castings.  It  can  also  readily  be  worked 
in  a  lathe.  German  silver  consists  of  about  80  to  95  per  cent 
brass  plus  5  to  10  per  cent  nickel. 

Bronzes  are  alloys  of  tin  and  copper.  Gun  metal  consists  of 
about  9  parts  copper  and  1  part  tin  ;  bell  metal  of  3  parts  copper 
and  1  part  tin.  Many  bronzes  also  contain  some  zinc  and  lead. 
This  is  particularly  the  case  with  bronze  used  for  statuary, 
which  consists  of  from  80  to  90  parts  copper,  3  to  8  parts  tin, 
1  to  3  parts  lead,  and  1  to  10  parts  zinc.  Bronze  used  in  ma- 
chinery commonly  contains  rather  more  lead  than  this,  and 
correspondingly  less  copper.  Phosphor  bronze  is  bronze  to  which 
from  0.5  to  3  per  cent  phosphorus  has  been  added.  It  is  particu- 
larly hard,  and  is  often  used  in  making  machinery.  Aluminum 
bronze  consists  of  from  5  to  10  per  cent  aluminum  and  90  to  95, 
per  cent  copper.  It  has  a  golden  yellow  color  and  is  quite  hard. 

Oxides  of  Copper.  —  Cuprous  oxide  Cu2O  is  obtained  as  one  of 
the  products  of  the  incomplete  oxidation  of  copper  when  the 
latter  is  slightly  heated  in  the  air.  By  reducing  hot,  alkaline 
solutions  of  cupric  salts  with  glucose,  cuprous  oxide  is  readily 
obtained;  also  by  heating  cuprous  chloride  with  sodium  car- 
bonate, thus : — 

2  CuCl  +  Na2CO3  =  Cu2O  +  CO2  +  2  NaCl. 

Cuprous  oxide  is  a  bright  red,  crystalline  powder  which  remains 
unchanged  in  the  air.  Dilute  acids  convert  it  into  cupric  salts 
and  copper,  thus :  — 

Cu20  +  2  HN03  =  Cu(N08)2  +  H2O  +  Cu. 
Cu2O  +  H2SO4  =  CuSO4  +  H2O  +  Cu. 

Cuprous  oxide  is  fused  with  glass  to  give  the  latter  a  red  color. 

Cupric  oxide  CuO  is  a  black  powder  formed  by  heating  copper 
in  oxygen  or  in  the  air.  It  is  also  obtained  by  ignition  of  the 
nitrate  or  carbonate.  It  readily  gives  up  its  oxygen  when  heated 
with  carbon  or  carbon  compounds  ;  hence  'the  use  of  cupric  oxide  in 
the  analysis  of  organic  compounds. 

Cuprous   hydroxide   is  not  known  with  certainty.      Cupric 


COPPER,    SILVER,    AND   GOLD  441 

hydroxide  Cu(OH)2  is  an  amorphous,  blue  precipitate  formed 
when  caustic  alkali  is  added  to  a  solution  of  a  cupric  salt.  The 
precipitate  readily  loses  water  on  boiling  it  in  the  solution  from 
which  it  has  separated,  thus  forming  a  black  hydrated  oxide  of 
approximately  the  composition  Cu(OH)2  •  2  CuO.  In  ammonia 
water  cupric  hydroxide  is  soluble,  yielding  the  deep  blue  solution 
called  Schweitzer's  reagent.  When  treated  with  acids,  cupric 
hydroxide  yields  cupric  salts. 

Halides  of  Copper.  —  Cuprous  chloride  CuCl  is  formed  by  boil- 
ing cupric  chloride  CuCl2,  with  metallic  copper  and  hydrochloric 
acid :  — 

CuCl2  +  Cu  =  2  CuCl. 

The  cuprous  chloride  separates  out  in  the  form  of  a  white,  finely 
crystalline  powder,  on  dilution  with  cold  water.  Cuprous  chlo- 
ride is  insoluble  in  water  and  also  in  alcohol;  but  it  dissolves 
in  hydrochloric  acid  or  ammonia.  The  hydrochloric  acid  solution 
absorbs  carbon  monoxide,  forming  a  product  which  separates  from 
the  concentrated  solutions  in  unstable,  shining  scales  that  prob- 
ably have  the  composition  CO  •  Cu2Cl2  •  2  H2O.  With  ammonia, 
cuprous  chloride  forms  the  compound  Cu2Cl2  •  2  NH3.  This 
ammoniacal  solution  also  has  the  property  of  absorbing  gases. 
In  many  ways,  cuprous  chloride  is  analogous  to  mercurous  chloride 
HgCl,  silver  chloride  AgCl,  and  thallous  chloride  T1CL 

Cupric  chloride  CuCl2  is  made  by  burning  copper  in  chlorine, 
also  by  dissolving  the  metal  in  nitro-hydrochloric  acid,  or  by 
treating  cupric  oxide,  hydroxide,  or  carbonate,  with  hydro- 
chloric acid.  The  anhydrous  salt  is  brown.  The  aqueous 
solutions  are  bluish  green,  and  from  them  green,  rhombic,  pris- 
matic crystals,  CuCl2  •  2  H2O,  may  be  obtained.  These  readily 
dissolve  in  water  or  in  alcohol.  At  500°  anhydrous  cupric 
chloride  melts,  and  on  further  ignition  it  may  be  converted 
into  cuprous  chloride.  With  ammonia,  cupric  chloride  forms 
CuCl2-6NH3,  a  dark  blue,  unstable  powder.  From  aqueous 
ammoniacal  solutions  of  cupric  chloride,  dark  blue  crystals  of 
the  composition  CuCl2  •  4  NH3  -  H2O  may  be  obtained.  At 
about  150°  these  compounds  are  converted  into  CuCl2  •  2  NH3, 
which  is  a  green  powder. 

Cuprous  bromide  CuBr,  like  cuprous  chloride,  is  a  white,  in- 
soluble powder.  It  may  be  melted  without  decomposition. 
Cuprous  iodide  Cul  is  the  only  iodide  of  copper  known.  It  is 


442  OUTLINES  OF   CHEMISTRY 

formed  by  heating  copper  and  iodine  together,  also  by  adding 
potassium  iodide  to  solutions  of  cupric  salts:  — 

CuSO4  +  2  KI  =  K2SO4  +  Cul  + 1. 

The  cuprous  iodide  so  formed  is  thus  precipitated  together 
with  free  iodine.  Cuprous  iodide  is  a  white,  insoluble,  crys- 
talline powder.  It  may  be  melted  at  red  heat.  Cuprous  fluo- 
ride CuF  is  also  known.  It  is  a  bright  red,  insoluble  powder. 

Cupric  bromide  CuBr2,  like  cupric  chloride,  is  readily  soluble 
in  water,  yielding  green  solutions.  Cupric  fluoride  CuF2  is 
also  known,  but  cupric  iodide  CuI2  is  not  known. 

Cyanides  of  Copper.  —  When  a  solution  of  a  cupric  salt  is 
treated  with  potassium  cyanide,  cupric  cyanide  Cu(CN)2  is 
formed  as  a  yellow  unstable  compound  which  at  once  decom- 
poses into  cuprous  cupric  cyanide  and  cyanogen,  thus  :  — 

CuS04  +  2  KCN  =  Cu(CN)2  +  K2SO4,  and 
3  Cu(CN)2  =  Cu(CN)2  .  Cu2(CN)2  +  (CN)2. 

This  reaction  is  frequently  used  in  making  cyanogen.  Cuprous 
cyanide  CuCN  may  be  obtained  as  a  white  precipitate  by 
adding  potassium  cyanide  to  a  solution  of  copper  sulphate 
saturated  with  sulphur  dioxide.  Potassium  cyanide  forms  a 
v  soluble  double  cyanide  with  cuprous  cyanide  CuCN  •  KCN. 
When  potassium  ferrocyanide  is  added  to  a  solution  of  a  cupric 
salt,  copper  ferrocyanide  forms  as  a  dark  brownish  red,  hydrous, 
amorphous  precipitate  :  — 

2  CuSO4  +  K4Fe(CN)6  =  Cu2Fe(CN)6  +  2  K2SO4. 

Copper  Salts  of  Oxy-acids.  —  No  cuprous  salts  of  oxy- acids  are 
known;  in  its  salts  with  such  acids,  copper  is  always  bivalent. 
The  following  are  the  most  important  of  these  salts :  — 

Copper  sulphate  CuSO4  is  obtained  by  heating  copper  with 
concentrated  sulphuric  acid.  On  a  commercial  scale  it  is 
formed  by  heating  copper  pyrites,  or  copper  plus  sulphur,  in 
a  current  of  air,  and  lixiviating  the  mass  with  water  and  evap- 
orating the  clear  solution,  from  which  large,  blue,  triclinic 
crystals  CuSO4  •  5  H2O,  known  as  blue  vitriol,  are  obtained 
(Fig.  75).  This  is  a  very  common  copper  salt.  On  exposure 
to  the  air  the  crystals  effloresce.  At  100°  they  lose  about 
four  fifths  of  their  water,  but  it  requires  about  200°  to 
dehydrate  them  completely.  The  anhydrous  salt  is  a  gray- 


COPPER,    SILVER,   AND   GOLD  443 

ish  white  powder,  insoluble  in  alcohol.  It  is  hygroscopic,  and 
is  consequently  sometimes  used  in  the  laboratory  as  a  drying 
agent.  Blue  vitriol  is  soluble  in  about  3  parts  of  water.  It  is 
used  in  large  quantities  in  copperplating,  as  a  mordant  in  dye- 
ing fabrics,  as  a  source  for  the  preparation  of  other  copper  com- 
pounds, and  as  a  germicide,  particularly  in  spraying  plants  to 
protect  them  from  insect  pests.  Mixed  with  calcium  hydroxide 
solution,  copper  sulphate  forms  the  so-called  Bordeaux  mixture, 
which  is  often  used  in  spraying  fruit  trees  and  other  plants. 

With  alkali  sulphates,  copper  sulphate  forms  double  salts,  like 
K2SO4  •  CuSO4  •  6  H2O.  These  are  monoclinic  and  isomorphous 
with  the  analogous  magnesium  salts.  With  ammonia,  copper 
sulphate  yields  deep  blue  solutions  from  which  dark  blue,  ortho- 
rhombic  crystals  of  the  composition  CuSO4  •  4  NH3  •  H2O  are 
precipitated  by  adding  alcohol.  On  heating  this  compound 
carefully,  it  may  be  converted  into  CuSO4  •  2  NH3,  a  green  pow- 
der which  is  called  cuprammonium  sulphate.  It  is  regarded  as 
ammonium  sulphate  in  which  two  hydrogen  atoms  are  replaced 
by  a  bivalent  copper  atom,  thus  :  — 

/NH8 

S04<^         p>Cu. 
NH3 

Basic  copper  sulphates,  of  the  formulae  CuSO4  •  Cu(OH)2, 
CuSO4  -  2  Cu(OH)2,  and  CuSO4  -  3  Cu(OH)2,  have  also  been 
prepared. 

Copper  nitrate  Cu(NO8)2  is  formed  by  dissolving  copper, 
cupric  oxide,  hydroxide,  or  carbonate  in  nitric  acid.  From 
the  solutions  the  salt  may  be  obtained  in  deliquescent  crystals, 
Cu(NO3)2  •  6  H2O.  On  heating  the  nitrate,  it  yields  cupric 
oxide. 

Copper  Carbonates. —  Only  basic  carbonates  of  copper  are  known. 
On  adding  an  alkaline  carbonate  to  a  solution  of  a  copper  salt, 
a  green  basic  copper  carbonate  CuCO3  -  Cu(OH)2  is  formed, 
thus : — 

2  CuS04  +  2  NagCOg  +  H2O 

=  2  Na2S04  +  CuCO3  •  Cu(OH)2  f  CO2. 

This  behavior  is  then  analogous  to  that  of  zinc  and  mag- 
nesium salts  when  they  are  similarly  treated  with  alkaline  car- 
bonates. Malachite,  which,  as  already  stated,  occurs  in  nature, 


444  OUTLINES   OF   CHEMISTRY 

is  CuCOg- Cu(OH)2,  and  verdigris,  which  forms  in  the  air 
on  copper  roofs,  etc.,  is  also  basic  copper  carbonate.  The 
mineral  azurite  forms  monoclinic  crystals  of  the  formula, 
Cu(OH)2.2CuCO3. 

Copper  Arsenite. —  Cu3(AsO3)2  •  2  H2O  is  a  beautiful,  green 
precipitate  formed  by  adding  potassium  arsenite  to  a  solution 
of  copper  sulphate.  It  is  known  as  Scheele's  green  and  is  used 
as  a  pigment. 

Sulphides  of  Copper.  —  Cuprous  sulphide  Cu2S  is  formed  by 
heating  cupric  sulphide  or  a  mixture  of  copper  and  sulphur  in 
a  current  of  l^drogen,  also  by  burning  copper  in  sulphur 
vapors.  It  forms  black  crystals  of  the  regular  system.  Cupric 
sulphide  is  obtained  as  a  black  precipitate  by  passing  hydrogen 
sulphide  into  a  solution  of  a  copper  salt.  It  is  practically  in- 
soluble in  water  and  also  in  dilute  acids. 

Analytical  Tests  for  Copper.  —  In  testing  copper,  the  following 
characteristic  reactions  are  generally  used:  — 

An  excess  of  ammonia  yields  deep  blue  solutions  with  copper 
salts.  Potassium  or  sodium  hydroxide  precipitates  blue  cupric 
hydroxide,  which  turns  black  on  boiling.  Hydrogen  sulphide 
produces  a  black  precipitate,  insoluble  in  dilute  acids.  Potas- 
sium ferrocyanide  produces  a  reddish  brown  precipitate  of 
copper  ferrocyanide.  Alkaline  carbonates  precipitate  basic 
copper  carbonate.  Hydrogen  sulphide  does  not  precipitate 
copper  sulphide  from  solutions  of  potassium  cuprous  cyanide; 
whereas,  in  a  similar  solution  of  potassium  cadmium  cyanide, 
yellow  cadmium  sulphide  is  precipitated,  and  thus  copper 
may  be  separated  from  cadmium. 

Occurrence,  Metallurgy,  and  Properties  of  Silver. — In  nature, 
silver  frequently  occurs  in  the  uncombined  condition,  usually 
scattered  in  quartz  and  other  rocks  in  small  amounts,  though 
at  times  it  is  found  in  masses  of  a  hundred  pounds  or  more. 
Native  silver  generally  is  contaminated  with  other  metals,  like 
gold,  copper,  iron,  and  lead.  It  is  frequently  found  in  fairly 
well-developed  crystals  of  the  regular  system.  The  most 
important  ores  of  silver  are  argentite  or  silver  glance  Ag2S, 
pyrargyrite  or  ruby  silver  ore  Sb2S3  •  3  Ag2S,  proustite  or  light 
red  silver  ore  As2S3  •  3  Ag2S,  and  other  complex  sulphides  of 
silver,  copper,  arsenic,  and  antimony.  Horn  silver  AgCl  and 
other  halides  of  silver  are  occasionally  found.  Nearly  all 


COPPER,   SILVER,   AND   GOLD  445 

galenite  PbS  contains  some  silver,  and  a  considerable  amount 
of  silver  is  obtained  from  this  source.  The  world's  supply  of 
silver  comes  mainly  from  the  United  States,  Mexico,  western 
South  America,  Australia,  Germany,  and  Austria. 

The  extraction  of  silver  from  its  ores  is  carried  on  by  several 
processes;  these  vary  according  to  the  nature  of  the  ore  and  the 
locality.  The  method  of  obtaining  silver  from  ores  rich  in  lead 
consists  of  roasting  the  ores  to  get  rid  of  the  sulphur,  and  then 
heating  the  residues  with  carbon  and  suitable  fluxes  in  a  blast 
furnace.  Thus  the  lead  containing  the  silver  in  solution 
settles  to  the  bottom  of  the  furnace  and  is  drawn  off.  To  this 
molten  lead,  zinc  is  added,  and  the  mixture  is  thoroughly 
stirred.  Zinc  is  but  slightly  soluble  in  lead,  but  silver  is  much 
more  greedily  taken  up  by  zinc  than  by  lead.  Therefore  the 
silver  enters  the  molten  zinc,  and  this  solution  of  silver  in  zinc 
floats  on  top,  where  it  solidifies  in  the  form  of  a  foam,  which  is 
skimmed  off  and  then  treated  with  superheated  steam  at  red 
heat.  This  oxidizes  the  zinc  and  leaves  the  silver  in  metallic 
form.  The  process  of  thus  using  zinc  in  separating  silver  from 
its  solutions  in  molten  lead  is  known  as  Parke's  process.  The 
amalgamation  process  consists  of  extracting  silver  from  its  ores  by 
dissolving  the  metal  in  mercury  and  then  distilling  off  the  latter. 
In  order  to  accomplish  this,  the  silver  in  the  ores  must  first  be 
converted  into  the  metallic  form.  This  is  done  by  roasting 
the  ores  with  sodium  chloride,  thus  forming  silver  chloride, 
which  is  then  moistened  with  water  and  treated  with  iron.  In 
this  way  ferrous  chloride  is  formed,  which  remains  dissolved, 
and  metallic  silver  is  precipitated.  The  latter  is  then  extracted 
from  the  mass  by  means  of  mercury.  Where  fuel  is  expensive, 
as  in  some  of  the  South  American  countries,  the  roasting  pro- 
cess is  not  used,  the  ores  first  being  finely  ground  and  then 
mixed  with  sodium  chloride,  copper  sulphate,  water,  and  mer- 
cury. This  thick  pasty  mass  is  mixed  by  driving  mules 
through  it,  and  shoveling  it  over  occasionally.  The  sodium 
chloride  and  copper  sulphate  react,  forming  cupric  chloride, 
which  in  turn  converts  the  silver  in  the  ores  to  silver  chloride. 
The  latter  is  then  decomposed  by  mercury,  forming  mercurous 
chloride  and  silver,  which  is  dissolved  in  the  excess  of  mercury 
present.  This  process  usually  requires  a  week  or  more.  The 
silver  amalgam  formed  is  washed  and  strained  through  canvas. 


446  OUTLINES   OF   CHEMISTRY 

From  the  solid  thus  remaining,  the  silver  is  obtained  by  dis- 
tilling off  the  mercury.  The  reactions  involved  are  in  the 
main  as  follows :  — 

CuSO4  +  2  NaCl  =Na2SO4  +  CuCl2. 

CuCl2  +  Ag2S  =  CuS  +  2  AgCl. 
2  AgCl  +  2  Hg  =  2  HgCl  +  2  Ag. 

Still  other  processes  are  also  sometimes  employed.  So,  when 
lead  containing  but  little  silver  cools,  the  crystals  that  form  are 
practically  pure  lead,  and  when  these  are  removed,  the  remain- 
ing liquid  is  richer  in  silver.  As  this  liquid  cools  further,  more 
solid  lead  forms,  after  the  removal  of  which  the  remaining  solu- 
tion is  still  richer  in  silver.  By  repeating  this  process,  known 
as  Pattinson's  process,  an  alloy  of  lead  relatively  rich  in  silver 
remains  behind,  from  which  the  silver  is  obtained  by  cupella- 
tion.  This  process  consists  of  heating  the  lead-silver  alloy  in 
a  current  of  air  in  a  muffle  furnace ;  the  lead  is  thus  oxidized 
to  lead  oxide,  which  melts  and  runs  off  or  is  volatilized,  and 
the  silver  remains  behind. 

Silver  is  refined  by  electrolysis  or  by  the  sulphuric  acid  method. 
The  latter  consists  of  dissolving  the  silver  in  boiling  hot,  con- 
centrated sulphuric  acid  and  then  diluting  the  solution  with 
water.  In  this  way,  the  noble  metals,  like  gold,  platinum, 
palladium,  etc.,  remain  behind,  and  the  silver  is  obtained  as 
silver  sulphate  in  solution.  From  the  latter  the  silver  is  pre- 
cipitated by  means  of  iron  and  then  cast  into  molds.  The 
electrolytic  process  of  refining  silver  is  similar  to  that  of  pre- 
paring pure  copper  by  electrolysis.  The  impure  silver  is  used 
as  anodes.  A  solution  of  silver  nitrate  acidified  with  nitric 
acid  serves  as  the  electrolyte,  and  thin  silver  plates  are  used 
as  cathodes,  on  which  silver  is  deposited  in  the  form  of  shining 
crystals  that  are  continually  rubbed  off  by  a  mechanical  con- 
trivance. These  crystals  of  pure  silver  fall  to  the  bottom  of 
the  vat,  and  after  draining  off  the  electrolyte  and  washing  them, 
they  are  ready  for  further  use.  The  electrolytic  process  of 
refining  silver  is  much  used  at  present.  In  the  chemical  lab- 
oratory pure  silver  is  often  made  by  heating  pure  silver  chloride 
with  soda,  or  by  treating  moist  silver  chloride  with  zinc. 

In  1910  the  United  States  produced  1958  tons  of  silver, 
which  is  nearly  a  third  of  the  world's  annual  production. 


COPPER,   SILVER,   AND   GOLD  447 

Silver,  argentum,  is  the  most  abundant  of  the  noble  metals. 
It  has  been  known  to  man  since  earliest  times,  and  has  been 
used  for  utensils  and  coins  for  many  centuries.  It  is  a  pure 
white  metal  of  high  metallic  luster.  Its  specific  gravity  is  10.6. 
At  962°  it  melts,  and  at  2050°  it  boils  and  may  be  distilled. 
Silver  is  very  tough,  ductile,  and  malleable,  and  may  be  beaten 
into  very  thin  sheets,  which  transmit  blue  light.  It  is  the  best 
conductor  of  heat  and  electricity  known.  Molten  silver  absorbs 
about  20  times  its  volume  of  oxygen,  which  is  again  liberated 
when  the  metal  cools,  thus  causing  "  spitting."  Silver  is  softer 
than  copper,  but  harder  than  gold.  Copper  is  often  added  to 
silver  to  secure  greater  hardness  and  durability.  Sterling  silver, 
which  is  used  for  coins,  spoons,  dishes,  and  many  other  useful 
articles,  is  an  alloy  containing  90  per  cent  silver  and  10  per  cent 
copper.  It  is  said  to  be  900  fine,  pure  silver  being  called  fine 
silver,  or  silver  that  has  a  fineness  of  1000.  Silver  also  readily 
forms  alloys  with  zinc,  lead,  mercury,  tin,  gold,  and  other 
metals.  When  alloyed  with  tin  and  mercury  it  forms  amal- 
gams which  are  used  for  filling  teeth. 

In  the  air,  silver  remains  unchanged,  but  ozone  oxidizes  it 
to  peroxide  AgO.  Molten  nitrates  or  hydroxides  of  the  alka- 
lies do  not  attack  silver,  hence  silver  dishes  are  used  in  the 
laboratory  in  working  with  caustic  alkalies  or  molten  saltpeter. 
Hydrochloric  acid  attacks  silver  but  slightly,  a  protective 
coating  of  silver  chloride  being  formed.  In  hot  concentrated 
sulphuric  acid,  silver  readily  dissolves,  liberating  sulphur 
dioxide :  — 

2  Ag  +  2  H2S04  =  Ag2S04  +  2  H2O  +  SO2. 

In  nitric  acid,  silver  is  readily  soluble  even  in  the  cold :  — 

3  Ag  +  4  HNO3  =  3  AgNO3  +  2  H2O  +  NO. 

Nitric  acid  is  the  best  solvent  for  silver.  With  sulphur,  silver 
readily  forms  silver  sulphide  Ag2S,  which  is  a  black  powder, 
insoluble  in  dilute  acids.  Silver  sulphide  forms  as  a  black  or 
dark  brown  coating  on  silver  spoons  when  in  contact  with  the 
contents  of  eggs,  for  the  latter  contain  sulphur.  With  the 
halogens,  silver  readily  unites  even  at  ordinary  temperatures. 

Oxides  of  Silver.  —  Silver  monoxide,  commonly  called  simply 
silver  oxide,  Ag2O,  is  formed  as  an  amorphous,  dark  brown  pre- 
cipitate when  potassium  or  sodium  hydroxide  is  added  to  a  silver 


448  OUTLINES   OF  CHEMISTRY 

nitrate  solution.  Above  250°,  silver  oxide  decomposes  into  the 
metal  and  oxygen,  a  behavior  which  is  similar  to  that  of  mer- 
curic oxide.  Silver  peroxide  AgO  forms  as  a  black  compound 
when  ozone  acts  on  silver.  It  is  also  formed  on  a  platinum 
anode  when  silver  nitrate  solutions  are  electrolyzed.  Silver 
suboxide  Ag4O  has  also  been  prepared.  Silver  hydroxide  is 
not  known. 

Halides  of  Silver.  — These  are  formed  by  the  action  of  halo- 
gens on  silver,  also  by  adding  a  halide  to  a  solution  of  a  silver 
salt,  or  by  treating  silver  oxide  with  hydrohalogen  acids. 

Silver  chloride  is  formed  as  a  curdy,  white  precipitate  when 
hydrochloric  acid  or  another  soluble  chloride  is  added  to  a 
solution  of  silver  nitrate.  It  is  insoluble  in  water  and  nitric 
acid,  but  it  dissolves  in  concentrated  hydrochloric  acid,  also  in 
ammonia  water,  alkaline  chlorides,  sodium  thiosulphate,  or 
potassium  cyanide.  Silver  chloride  is  white,  but  on  exposure 
to  light  it  turns  violet  and  then  brown  or  black,  probably 
because  of  the  separation  of  finely  divided  silver.  The  salt 
melts  at  151°.  It  is  much  used  in  analytical  chemistry  in  the 
detection  and  estimation  of  silver  and  also  of  chlorine.  Silver 
chloride  readily  absorbs  ammonia,  forming  AgCl-3NH3. 

Silver  bromide  AgBr  is  obtained  as  a  slightly  yellowish  pre- 
cipitate by  adding  a  soluble  bromide  to  a  silver  nitrate  solution. 
If  a  soluble  iodide  is  added  instead  of  the  bromide,  silver  iodide 
Agl,  a  yellow  precipitate,  is  produced.  Both  the  bromide  and 
the  iodide  of  silver  are  also  practically  insoluble  in  water  and 
dilute  acids.  The  solubility  of  silver  bromide  in  ammonia  is 
less  than  that  of  the  chloride;  and  the  iodide  is  least  soluble  in 
ammonia.  When  exposed  to  light,  the  bromide  and  iodide  of 
silver  are  partially  decomposed,  like  the  chloride.  Silver  fluo- 
ride AgF  is  very  soluble  in  water,  and  consequently  differs 
greatly  from  the  other  halides  of  silver.  It  may  be  obtained 
by  adding  hydrofluoric  acid  to  silver  oxide.  The  crystals  are 
deliquescent  and  correspond  to  the  formula  AgF  •  H2O  or 
AgF  •  2  H20. 

Use  of  Silver  Halides  in  Photography.  —  The  fact  that  the 
halides  of  silver  are  decomposed  when  exposed  to  light  forms 
the  basis  of  modern  photography.  The  sensitive  dry  plates  or 
films  are  covered  with  a  layer  of  gelatine  containing  finely 
divided  silver  bromide  in  suspension.  When  such  a  plate  is 


COPPER,    SILVER,    AND   GOLD  449 

exposed  to  light  in  a  camera,  the  silver  bromide  is  acted  upon 
more  strongly  in  some  places  than  in  others,  according  to  the 
intensity  of  the  light  reflected  from  the  different  parts  of  the 
object  whose  image  is  focused  upon  the  plate.  After  being 
thus  momentarily  exposed  to  light,  the  plate  exhibits  no  visible 
change.  However,  on  now  treating  the  plate  with  reducing 
solutions  like  ferrous  salts,  alkaline  pyrogallic  acid,  or  hydro- 
chinone,  silver  is  deposited  from  the  bromide  more  rapidly  and 
densely  on  those  places  where  the  light  has  acted  more  intensely, 
and  less  rapidly  where  the  action  of  the  light  has  been  less  in- 
tense. Thus  these  reducing  solutions,  or  developers,  as  they 
are  called,  develop  the  picture,  which  is  a  so-called  negative, 
for  it  is  dark  where  the  object  was  light,  and  light  where  the 
object  was  dark.  The  developing  process  must  be  stopped 
when  a  good  picture  of  sharply  denned  outline  has  been  secured. 
This  is  done  by  rinsing  the  developing  solution  from  the  plate 
and  introducing  the  latter  into  a  solution  of  sodium  thiosul- 
phate  Na2S2O3,  commonly  called  hyposulphite  of  soda  or 
"hypo."  Thus  the  developing  process  is  arrested,  and  the 
sodium  thiosulphate  dissolves  the  remaining  unaltered  silver 
bromide  from  the  plate,  forming  a  soluble  sodium  silver  thio- 
sulphate and  sodium  bromide,  which  are  finally  rinsed  off :  — 

AgBr  +  Na2S2O3  =  NaBr  +  NaAgS2O3. 

The  film  or  the  plate  now  consists  simply  of  silver  imbedded 
in  gelatine  and  is  said  to  be  "  fixed,  "  for  it  may  be  exposed  to 
the  light  without  suffering  further  alteration  thereby.  The 
sodium  thiosulphate  solution  is  termed  the  "fixing  bath." 
The  pictures  are  printed  by  placing  the  negative  upon  paper 
covered  with  a  sensitive  film  of  silver  chloride,  bromide,  or 
iodide,  and  then  exposing  this  to  the  light.  This  yields  a 
positive,  for  obviously  the  print  is  dark  where  the  negative  is 
light  and  vice  versa.  The  prints  must  be  "  fixed  "  by  the  same 
process  as  the  negative.  Usually  silver  bromide  papers  are 
employed  for  printing.  They  are  more  readily  acted  upon  by 
light  than  papers  covered  with  a  silver  chloride  film. 

Silver  Nitrate  AgNO3,  also  known  as  argentic  nitrate  and 
lunar  caustic,  is  the  most  important  silver  compound,  for  from  it 
nearly  all  the  other  compounds  are  directly  or  indirectly  prepared. 
Silver  nitrate  is  formed  by  dissolving  silver  in  nitric  acid. 

2G 


450  OUTLINES   OF  CHEMISTRY 

It  forms  transparent,  rhombic  crystals  which  melt  at  about 
200°,  and  are  extremely  soluble  in  water,  although  they  are  an- 
hydrous and  not  at  all  deliquescent.  At  0°,  100  parts  of  water 
dissolve  122  parts  of  silver  nitrate,  while  at  100°  nearly  ten 
times  that  amount  of  salt  is  dissolved.  In  contact  with  organic 
matter,  like  the  skin,  cloth,  or  paper,  silver  nitrate  turns  black 
on  exposure  to  light,  owing  to  separation  of  metallic  silver. 
Hence  the  salt  is  used  in  indelible  inks.  Silver  nitrate  has 
strong  caustic  properties,  and  hence  is  used  in  cauterizing 
sores  and  removing  warts.  For  this  purpose  it  is  commonly 
cast  into  sticks,  which  are  either  pure  silver  nitrate  or  a  mixture 
of  the  latter  with  potassium  nitrate.  It  is  in  the  form  of  such 
caustic  pencils  that  silver  nitrate  is  termed  lunar  caustic  or  lapis 
infernis.  Silver  nitrate  is  frequently  used  in  analytical  chem- 
istry. It  is  poisonous,  and  its  solutions  have  a  disagreeable 
metallic  taste.  With  ammonia,  silver  nitrate  forms  rhombic 
crystals  of  the  composition  — 

AgN03.2NH3. 

Silver  Nitrite  AgNO2  is  obtained  as  a  sparingly  soluble 
crystalline  precipitate  by  adding  potassium  nitrite  to  a  solution 
of  silver  nitrate.  The  salt  is  used  in  water  analysis. 

Silver  Sulphate  Ag2SO4  is  commonly  made  by  dissolving 
silver  or  silver  carbonate  in  concentrated  sulphuric  acid.  It 
forms  rhombic  crystals  soluble  in  about  200  parts  of  water. 
In  sulphuric  acid  it  dissolves  more  readily,  because  of  the  for- 
mation of  silver  bisulphate  AgHSO4. 

Silver  Carbonate  Ag2CO3  is  obtained  as  a  yellowish  precipi- 
tate by  adding  soluble  carbonates  to  solutions  of  silver  salts. 
It  dissolves  in  water  charged  with  carbon  dioxide.  On  heating 
silver  carbonate,  it  decomposes  :  — 


Silver  Phosphate  Ag3PO4  is  a  yellow,  amorphous  precipi- 
tate produced  by  adding  sodium  phosphate  to  a  solution  of 
a  silver  salt. 

Silver  Sulphide  Ag2S  is  obtained  as  a  dark  brown  precipi- 
tate, when  hydrogen  sulphide  acts  on  solutions  of  silver  salts. 
It  is  insoluble  in  dilute  acids. 

Silver  Cyanide  AgCN  is  formed  by  adding  potassium  cya- 
nide to  a  solution  of  silver  nitrate.  In  excess  of  potassium 


COPPER,   SILVER,   AND   GOLD  451 

cyanide,  silver  nitrate  is  soluble,  forming  the  double  cyanide 
KAg(CN)2.  The  latter  is  very  important,  for  its  solutions  are 
used  in  silver  plating.  Silver  cyanide  is  stable  in  the  light, 
and  is  used  in  pharmaceutical  practice  for  preparing  dilute 
solutions  of  hydrocyanic  acid  by  the  following  reaction  : — 

AgCN  +  HCl  =  AgCl  +  HCN. 

The  silver  chloride  being  insoluble  is  filtered  off. 

Silver  Plating.  —  This  is  commonly  accomplished  by  elec- 
trolysis. The  objects  to  be  plated  are  first  thoroughly  cleaned, 
and  then  immersed  in  a  bath  consisting  of  a  solution  of  potas- 
sium silver  cyanide  KAg(CN)2,  in  which  they  serve  as 
cathodes.  The  anode  consists  of  a  thick  plate  of  silver.  As 
the  current  passes,  the  objects  are  coated  with  a  dense,  white, 
well-adhering  deposit  of  silver  which  can  afterwards  be  pol- 
ished. Aqueous  silver  nitrate  will  not  serve  as  the  electrolyte  in 
electroplating,  for  from  such  solutions  crystalline,  poorly  adhering 
deposits  are  obtained. 

Silver  mirrors  are  formed  when  a  clean  surface  of  glass  is 
treated  with  an  ammoniacal  silver  nitrate  solution  plus  a 
reducing  agent  like  formic  aldehyde,  glucose,  or  Rochelle  salt. 

Silver  Fulminate  AgONC  is  made  like  mercuric  fulminate. 
It  is  even  more  explosive  than  the  latter. 

Analytical  Tests  for  Silver.  —  Silver  is  very  readily  detected 
in  its  compounds.  All  silver  salts  of  organic  acids  yield  white 
metallic  silver  when  ignited.  Silver  is  precipitated  from  solu- 
tions of  silver  salts  by  copper,  mercury,  zinc,  or  iron.  Solu- 
tions of  silver  salts  yield  insoluble  silver  chloride  on  treatment 
with  hydrochloric  acid.  Silver  chloride  is  soluble  in  ammonia, 
which  is  not  the  case  with  mercurous  chloride  or  lead  chloride. 
The  last  two  salts  are  also  insoluble  in  water,  though  lead 
chloride  dissolves  in  boiling  water.  Silver  chromate  Ag2CrO4, 
a  dark  red  precipitate,  forms  by  adding  a  solution  of  either  po- 
tassium chromate  or  bichromate  to  a  soluble  silver  salt.  Other 
characteristic  precipitates  are  the  carbonate,  which  dissolves  in 
ammonium  carbonate,  and  the  phosphate,  which  is  yellow ;  these 
have  already  been  described.  In  potassium  cyanide  or  in 
sodium  thiosulphate  solutions,  silver  salts  are  soluble.  The 
fact  that  most  silver  compounds  are  affected  by  light  is  also 
characteristic. 


452  OUTLINES   OF  CHEMISTRY 

Occurrence,  Metallurgy,  and  Properties  of  Gold.  —  Gold  is 
usually  found  in  nature  in  the  uncombined  state  in  quartz 
veins  and  alluvial  sands,  though  sometimes  it  occurs  combined 
with  tellurium  as  calaverite  AuTe2,  petzite  (AuAg)2Te,  and 
sylvanite  AgAuTe2.  Native  gold  always  contains  some  sil- 
ver, frequently  also  copper,  lead,  iron,  and  other  metals.  Many 
samples  of  iron  pyrites  contain  small  amounts  of  gold.  The 
chief  localities  are  South  Africa,  Alaska,  California,  Colorado, 
Dakota,  the  Urals,  and  Australia. 

Alluvial  gold-bearing  sands  are  usually  washed  with  water, 
which  carries  away  the  light,  loose  material,  leaving  the  heavier 
particles  of  gold  behind.  The  latter  is  gathered  by  treatment 
with  mercury;  and  from  the  gold  amalgam  thus  formed,  the 
mercury  is  removed  by  distillation  from  iron  retorts.  From 
quartz,  gold  is  separated  by  pulverizing  the  material  in  stamp- 
ing mills,  and  then  running  water  bearing  the  finely  divided 
ore  over  copper  plates  amalgamated  with  mercury.  The  latter 
thus  catches  and  dissolves  the  gold  particles.  From  time  to 
time,  the  amalgam  is  removed  and  distilled  to  separate  the 
gold  from  the  mercury.  The  latter  is  used  over  and  over 
again.  Not  all  the  gold  is  extracted  from  the  ores  by  this 
amalgamation  process,  the  "  tailings "  generally  containing 
very  finely  divided  particles  of  the  metal.  These  are  recovered 
by  the  cyanide  process,  which  consists  of  treating  the  ore  with 
a  dilute  solution  of  potassium  cyanide  (from  0.25  to  0.8  per 
cent),  in  which  both  gold  and  silver  dissolve  in  presence  of  the 
oxygen  of  the  air,  thus :  — 

4  Au  +  8  KCN  +  2  H2O  +  O2  =  4  KOH  +  4  KAu(CN)2. 
4  Ag  +  8  KCN  +  2  H2O  +  O2  =  4  KOH  +  4  KAg(CN)2. 

These  cyanide  solutions  are  then  subjected  to  electrolysis  to 
obtain  the  gold,  or  the  latter  is  precipitated  by  means  of  zinc, 
thus : — 

2  KAu(CN)2  +  Zn  =  K2Zn(CN)4  +  2  Au. 

If  electrolysis  is  used,  iron  anodes  and  lead  cathodes  are  em- 
ployed. The  latter  are  finally  melted  and  cupelled,  by  heating 
them  in  a  current  of  air  in  a  muffle  furnace.  Thus  lead  is 
transformed  to  oxide  and  runs  off,  and  gold  remains.  Aurife- 
rous pyrites  ores  are  roasted,  and  then  subjected  to  the 
chlorination  process,  which  consists  of  treating  the  finely  ground 


COPPER,   SILVER,   AND   GOLD  453 

ore  with  chlorine,  thus  forming  soluble  auric  chloride  AuCl3. 
From  the  solution,  gold  is  then  precipitated  by  means  of  ferrous 
sulphate,  thus :  — 

2  AuCl3  +  6  FeS04  =  2  FeCl3  +  2  Fe2(SO4)3  +  2  Au. 
Gold  from  any  of  the  above  processes  usually  contains  silver 
from  which  it  must  be  parted.  The  parting  is  accomplished 
by  treating  the  alloy  in  form  of  foil  or  granulated  pieces  with 
hot  concentrated  sulphuric  acid  which  dissolves  the  silver  and 
leaves  the  gold  behind  as  a  dark  powder.  This  is  finally 
melted  with  fluxes,  like  potassium  nitrate  or  borax,  to  remove 
any  further  base  metals  that  may  still  be  present.  Gold  may 
also  be  separated  from  silver  by  electrolytic  methods.  In  the 
process  of  assaying  gold  ores,  the  latter  are  heated  with  lead 
and  fluxes  like  borax  or  soda.  Thus  gold  and  silver  finally 
accumulate,  dissolved  in  lead.  This  lead  alloy,  or  "  button,"  is 
then  heated  in  a  muffle  furnace  in  a  small  dish  called  a  cupel, 
made  of  bone  ash,  which  absorbs  the  lead  oxide  formed,  leaving 
the  alloy  of  gold  and  silver.  This  is  then  treated  with  hot 
nitric  acid  which  dissolves  the  silver  and  leaves  the  gold 
behind.  This  method  of  parting  is  called  quartation,  for  it 
was  formerly  believed  that  in  order  that  the  separation  be  com- 
plete, the  alloy  must  not  contain  more  than  25  per  cent  gold. 
As  a  matter  of  fact,  the  alloy  may  contain  up  to  35  per  cent 
gold  and  still  the  separation  be  successful. 

Pure  gold  is  a  rather  soft,  bright  yellow  metal  having  a  high 
luster.  It  is  the  most  malleable  and  ductile  of  all  the  metals,  and 
conducts  heat  and  electricity  we41.  It  may  be  beaten  into  very 
thin  leaves,  which  transmit  green  light.  A  grain  of  gold  may 
be  beaten  into  a  leaf  covering  an  area  of  about  half  a  square 
meter.  Gold  may  be  obtained  in  regular  octahedra  or  dodeka- 
hedra.  Its  specific  gravity  is  19.3.  It  melts  at  1064°,  and 
may  be  volatilized  in  the  electric  furnace.  Chemically,  gold  is 
rather  inert.  In  the  air  or  in  oxygen,  it  remains  unchanged  even 
at  high  temperatures.  Gold  is  attacked  by  fused  caustic  alkalies 
and  nitrates  of  the  alkalies,  aurates  being  formed.  It  is  soluble 
in  aqua  regia,  but  not  in  nitric,  hydrochloric,  or  sulphuric  acid. 
Chlorine  attacks  gold,  but  sulphur  does  not.  In  potassium 
cyanide,  it  readily  dissolves,  as  already  stated.  On  ignition, 
all  compounds  of  gold  are  decomposed,  leaving  a  residue  of  the 
metal.  The  atomic  weight  of  gold  is  197.2. 


454  OUTLINES   OF   CHEMISTRY 

Gold  Alloys.  —  Gold  forms  alloys  with  many  metals,  like  mer- 
cury, copper,  silver,  cadmium,  tin,  and  lead.  Pure  gold  is  too 
soft  for  coinage,  jewelry,  and  ornaments;  hence  it  is  alloyed  with 
copper  for  these  purposes.  Indeed,  of  the  gold  alloys  those  with 
copper  are  by  far  the  most  important  ones.  They  are  harder, 
stronger,  darker  in  color,  and  more  readily  fusible  than  pure 
gold.  The  gold  coins  of  the  United  States  consist  of  1  part 
copper  and  9  parts  gold.  It  is  customary  to  express  the  purity 
of  gold  in  carats.  Pure  gold  is  called  24-carat  gold ;  a  gold 
alloy  consisting  of  8  parts  copper  and  16  parts  gold  is  16-carat; 
14-carat  gold  contains  10  parts  copper  and  14  parts  gold.  For 
many  articles  of  jewelry,  it  is  necessary  to  use  at  least  as  low  as 
14-carat  gold  to  secure  sufficient  rigidity. 

Pure  gold  leaf  in  form  of  a  wad  is  readily  condensed  to  a 
solid  piece  by  hammering,  which  fact  is  used  by  dentists  in 
filling  teeth  with  gold. 

Gold  has  been  known  since  earliest  times  and  has  always 
"been  regarded  as  an  article  of  great  value.  In  1910  the  United 
States  produced  159.6  tons  of  gold,  valued  at  $96,269,100. 
This  is  approximately  one  fourth  of  the  world's  annual  output. 
The  yield  of  the  Transvaal  mines  is  now  greater  than  that  of 
America.  The  production  of  gold  has  increased  greatly  in 
recent  years.  Only  about  26  tons  of  gold  were  produced 
annually  on  the  average  during  the  first  fifty  years  of  the  last 
century.  Since  1850  there  has  been  a  steady  increase  in  the 
production  of  gold. 

Compounds  of  Gold.  —  On  account  of  its  inertness,  gold  forms 
but  few  compounds.  There  are  aurous  compounds  in  which 
gold  is  univalent,  and  auric  compounds  in  which  it  is  trivalent. 
When  dissolved  in  aqua  regia  or  when  treated  with  chlorine, 
auric  chloride  AuCl3  is  formed.  It  consists  of  reddish-brown 
deliquescent  crystals  which,  when  heated  to  about  180°, 
decompose  into  chlorine  and  aurous  chloride  AuCl,  which  is 
white.  On  boiling  aurous  chloride  with  water,  auric  chloride 
and  gold  are  formed :  — 

3  AuCl  =  AuCl3  +  2  Au. 

When  auric  chloride  is  evaporated  with  excess  of  hydrochloric 
acid,  yellow  prismatic  crystals  of  chlorauric  acid  HAuCl4-4  H2O 
are  formed.  With  sodium  chloride,  auric  chloride  forms 


COPPER,    SILVER,   AND   GOLD  455 

sodium  chloraurate  NaAuCl4  or  NaCl  •  AuCl3,  which  is  a  salt  of 
chlorauric  acid.  This  salt  is  used  in  photography  in  dilute  solu 
tions  as  a  toning  bath,  for  silver  prints  immersed  in  it  are  col- 
ored slightly  yellowish,  due  to  the  fact  that  silver  decomposes 
the  salt  and  so  deposits  a  very  thin  layer  of  gold.  The  analo- 
gous bromides  and  iodides  of  gold  are  also  known.  They  are 
rather  unstable. 

Aurous  oxide  Au2O  is  a  dark  violet  powder  formed  by  the 
action  of  caustic  alkalies  upon  aurous  chloride.  When  heated, 
it  yields  gold  and  oxygen.  Auric  oxide  Au2O3  is  a  brown 
powder.  Auric  hydroxide  Au(OH)3  is  formed  when  caustic 
alkalies  are  added  to  auric  chloride  solutions.  This  hydroxide 
is  soluble  in  excess  of  caustic  alkalies,  forming  aurates  like 
KAuO2  :  — 

Au(OH)3  +  KOH  =  KAu02  +  2  H2O. 

By  treating  auric  chloride  or  oxide  with  ammonia,  a  precipi- 
tate is  formed  which  is  called  fulminating  gold.  It  is  very 
explosive  when  dry.  Its  composition  is  not  known  with  cer- 
tainty. 

Aurous  cyanide  AuCN  is  a  yellow,  crystalline  powder  which 
is  soluble  in  potassium  cyanide,  forming  potassium  aurous  cya- 
nide KAu(CN)2.  The  latter  is  also  obtained  by  dissolving 
gold  in  dilute  potassium  cyanide  solutions  in  presence  of  air. 
When  auric  compounds  are  treated  with  potassium  cyanide, 
potassium  auric  cyanide  KAu(CN)4  is  formed.  The  solutions 
of  these  double  cyanides  of  potassium  and  gold  are  used  as 
baths  in  gold  electroplating,  which  process  is  similar  to  silver 
plating.  A  gold  plate  serves  as  the  anode,  and  the  object  to 
be  plated  is  the  cathode. 

Auric  sulphide  Au2S3  is  formed  as  brownish  black  precipitate 
when  hydrogen  sulphide  is  passed  into  cold  solutions  of  auric 
salts.  The  precipitate  usually  contains  free  sulphur.  Auric 
sulphide  is  soluble  in  alkali  sulphides,  sulpho-salts,  like 
(NH4)3AuS3,  being  formed.  From  hot  solutions,  aurous  sul- 
phide Au2S  is  similarly  precipitated.  It  has  a  steel-gray 
appearance  and  dissolves  in  water,  from  which  it  may  be  pre- 
cipitated by  adding  hydrochloric  acid. 

Purple  of  Cassius  consists  of  a  finely  divided,  brownish  pur- 
ple precipitate  of  gold  formed  by  adding  stannous  chloride  to 


456  OUTLINES   OF   CHEMISTRY 

solutions  of  auric  chloride.      The  powder  is  used  in  coloring 
glass  and  in  painting  porcelain  ware. 

Analytical  Tests  for  Gold.  —  On  ignition  with  soda  on  char- 
coal, gold  compounds  yield  a  globule  of  gold.  Ferrous  sulphate 
precipitates  gold  from  solutions  of  its  salts.  Other  reducing 
agents,  like  oxalic  acid  and  sulphur  dioxide,  also  precipitate 
gold  in  brown,  pulverulent  form.  Gold  is  readily  displaced 
from  its  solutions  by  many  other  metals  like  copper,  zinc,  and 
iron.  The  purple  of  Cassius  reaction  already  mentioned  is 
also  characteristic. 


CHAPTER   XXVI 

THE  METALS  OF  THE  EARTHS 

OF  the  metals  of  the  earths,  aluminum  is  by  far  the  most 
important.  It  has  already  been  stated  that  boron  is  a  trivalent 
element  and  that  its  compounds  consequently  bear  some  analogy 
to  those  of  aluminum.  Boron,  however,  exhibits  but  slightly 
basic  properties,  whereas  aluminum  is  a  basic  element,  though 
its  hydroxide  is  also  capable  of  acting  as  a  weak  acid  in  the 
formation  of  alum  mates. 

The  other  metals  of  the  earths  are  gallium,  indium,  thallium, 
scandium,  yttrium,  lanthanum,  and  ytterbium.  With  the 
exception  of  thallium,  these  are  all  quite  rare.  In  their  com- 
pounds, these  rare-earth  metals  are  trivalent  like  aluminum. 

Occurrence,  Preparation,  and  Properties  of  Aluminum.  — Alu- 
minum is  very  widely  distributed.  It  is  the  most  abundant  of 
the  metals  and  enters  into  the  composition  of  the  earth's  crust 
to  the  extent  of  nearly  8  per  cent.  Aluminum  has  a  great 
affinity  for  oxygen,  and  so  is  never  found  in  the  uncombined 
state  in  nature.  It  occurs  as  an  essential  ingredient  of  practi- 
cally all  the  common  siliceous  rocks.  It  is  found  particularly  in 
feldspars,  micas,  chlorite,  granites,  slates,  and  clays.  Oxide  of 
aluminum  occurs  as  corundum,  also  in  form  of  sapphire  or  ruby. 
Bauxite  is  a  mineral  consisting  of  the  hydrated  oxides  of  alumi- 
num and  iron.  Cryolite,  which  is  found  in  Greenland,  is  a  double 
fluoride  of  aluminum  and  sodium  of  the  formula  A1F3 .  3  NaF. 
Though  widely  distributed,  aluminum  is  not  found  in  animals, 
and  only  in  traces  in  some  plants. 

Aluminum  was  formerly  prepared  by  heating  sodium  alumi- 
num chloride  NaCl  •  A1C13,  or  cryolite,  with  metallic  sodium. 
It  is  now  made  in  large  quantities  by  electrolysis  of  a  solution 
of  aluminum  oxide  A12O3  in  molten  cryolite.  The  containing 
vessel  (Fig.  150)  is  made  of  graphitic  carbon,  which  also  serves 
as  cathode.  The  anode  consists  of  sticks  of  carbon  placed 
vertically.  Thus,  as  the  current  passes,  oxygen  is  evolved  on 

457 


458 


OUTLINES   OF  CHEMISTRY 


the  carbon  anode  and  passes  off.  Aluminum  is  deposited  at 
the  bottom  of  the  containing  vessel  and  is  tapped  off  from 
time  to  time.  The  heat  generated  by  the  current  depositing 
the  metal  is  sufficient  to  keep  the  electrolyte  and  the  aluminum 


FIG.  150. 

in  a  molten  state.  As  the  aluminum  is  deposited,  more  oxide 
is  placed  on  top  of  the  electrolyte,  as  shown  in  the  figure. 
Large  quantities  of  aluminum  are  now  prepared,  especially 
where  water  power  is  available,  as  at  Niagara  Falls  and  other 
localities.  In  1910  the  United  States  produced  about  24,000 
tons  of  aluminum. 

Aluminum  is  a  silver-white,  lustrous  metal  of  specific  gravity 
2.6  to  2.7;  that  is,  it  is  about  one  third  as  heavy  as  iron.  It 
melts  at  about  660°.  It  is  very  ductile  and  malleable  and  is 
a  good  conductor  of  heat  and  electricity.  Aluminum  is  about 
as  hard  as  silver.  The  hammered  or  rolled  metal  is  denser 
than  when  cast,  which  is  also  in  general  true  of  other  metals. 
Sufficiently  large  pieces  of  aluminum  ring  like  a  bell  when 
struck.  At  a  temperature  slightly  below  its  melting  point, 
aluminum  becomes  brittle  and  crumbles  when  shaken  or 
touched.  At  still  lower  temperatures,  it  is  again  pliable  and 
may  be  welded  and  readily  worked  into  desired  forms.  It  is 
very  difficult  to  solder  aluminum,  and  consequently  it  cannot 
very  well  be  used  for  many  purposes.  Chemically  aluminum 
is  relatively  inert  as  compared  with  the  metals  of  the  alkalies, 
alkaline  earths,  and  magnesium.  It  oxidizes  but  very  slowly 
on  exposure  to  the  air  or  oxygen ;  for  a  thin,  though  resistant, 
coating  of  oxide  forms  on  the  metal,  giving  it  a  slightly  bluish 


THE  METALS  OF  THE  EARTHS  459 

hue.  This  film  protects  the  aluminum  from  further  oxidation 
and  from  attack  by  many  acids.  Thin  pieces  of  aluminum, 
when  strongly  heated  in  the  air  or  in  oxygen,  burn  with  a  bril- 
liant light,  forming  the  oxide  and  some  nitride.  The  metal 
may  also  be  burned  in  a  current  of  superheated  steam,  oxide 
and  hydrogen  being  formed.  In  caustic  alkalies,  aluminum 
dissolves,  yielding  'aluminates  and  hydrogen.  Hydrochloric 
acid  readily  acts  on  aluminum,  forming  hydrogen  and  alumi- 
num chloride,  which  is  soluble.  Nitric  acid  and  dilute  sul- 
phuric acid  are  practically  without  action  on  aluminum,  which 
is,  however,  soluble  in  hot,  concentrated  sulphuric  acid  with 
concomitant  evolution  of  sulphur  dioxide.  The  atomic  weight 
of  aluminum  is  27.1;  the  metal  is  always  trivalent. 

Uses  of  Aluminum.  —  On  account  of  the  fact  that  aluminum 
is  a  good  conductor  of  electricity,  it  is  used  for  electric  cables 
and  wires,  especially  at  times  when  copper  is  relatively  high 
in  price.  Aluminum  is  used  for  cooking  utensils,  and  many 
other  useful  articles.  In  finely  divided  form,  it  is  used  in 
aluminum  paints.  In  the  form  of  leaf,  it  is  employed  in  stamp- 
ing titles  on  the  covers  of  books,  for  it  does  not  blacken  on 
exposure  to  the  air  as  silver  does.  Aluminum  is  further  em- 
ployed in  removing  oxides  from  iron,  and  thus  denser  castings 
are  obtained.  Aluminum  alloys  are  also  in  use.  Besides  the 
aluminum-bronze  already  mentioned,  magnalium,  consisting  of 
10  to  25  per  cent  magnesium  and  90  to  75  per  cent  aluminum, 
is  coming  into  use.  It  is  lighter  and  much  harder  than  alu- 
minum and  may  be  polished  to  a  higher  degree.  With  cadmium, 
aluminum  forms  a  very  tough  alloy.  Alloys  with  nickel,  zinc, 
and  other  metals  have  also  been  studied. 

On  account  of  the  great  affinity  which  aluminum  has  for 
oxygen,  it  is  used  in  preparing  other  metals  from  their  oxides 
by  the  Goldschmidt  process,  which  consists  of  heating  together  a 
mixture  of  powdered  aluminum  and  the  oxide  to  be  reduced. 
Thus  the  oxides  of  iron,  nickel,  chromium,  etc.,  are  readily 
reduced;  for  example, 

3  Fe3O4  -f  8  Al  =  4  A12O3  +  9  Fe. 
Fe203  +  2  Al  =  A1203  +  2  Fe. 
3N10  +  2  Al  =  Al208  +  3Ni. 
Cr2O3  +  2  Al  =  A12O3  +  2  Cr. 


460  OUTLINES   OF  CHEMISTRY 

In  practice  the  mixture  of  the  oxide  and  aluminum  powdei 
is  placed  in  a  crucible  and  ignited  by  means  of  a  luse  of  magne- 
sium ribbon  or  a  mixture  of  either  aluminum  or  magnesium  and 
barium  peroxide.  The  reaction,  when  once  started,  continues 
with  great  evolution  of  heat.  The  temperature  attained  during 
the  reaction  when  ferric  oxide  is  reduced  by  aluminum  is  about 
3000°,  which  is  quite  sufficient  to  melt  both  the  iron  and  the 
aluminum  oxide  formed.  Upon  this  fact  Goldschmidt  has  based 
his  famous  method  of  welding  iron,  which  consists  essentially  of 
butting  together  the  parts  to  be  welded,  surrounding  the  joint 
with  an  ignited  mixture  of  oxide  of  iron,  usually  Fe3O4,  and 
powdered  aluminum.  This  mixure  is  called  thermite.  The  heat 
developed  by  the  reaction  is  sufficient  to  weld  the  iron  securely. 
Thus  car  rails  can  be  welded  when  in  place,  and  many  repairs 
on  machinery,  ships,  etc.,  can  conveniently  be  made.  Thermite 
welding  is  consequently  frequently  used. 

Many  metallic  sulphides,  when  heated  with  aluminum  powder, 
may  similarly  be  reduced  to  metal,  aluminum  sulphide  being 
simultaneously  formed. 

Aluminum  Oxide  or  Alumina  A12O8  is  found  in  nature  as 
corundum,  ruby,  or  sapphire  in  hexagonal  crystals.  Corundum 
is  colorless,  ruby  is  red  due  to  the  presence  of  a  little  chromium, 
and  sapphire  is  blue  because  it  contains  a  trace  of  cobalt.  Im- 
pure corundum  is  called  emery.  It  contains  ferric  oxide,  and  is 
used  as  an  abrasive  material  on  account  of  its  great  hardness, 
which  is  but  slightly  below  that  of  the  diamond.  Aluminum 
oxide  results  in  the  Goldschmidt  reduction  process,  above  de- 
scribed, also  when  the  hydroxide  of  aluminum  is  strongly 
ignited.  The  oxide  thus  obtained  is  practically  not  attacked 
by  acids.  After  fusion  with  bisulphate  of  potassium  or 
with  caustic  alkalies,  it  may  be  dissolved  in  water,  for  thus 
aluminum  sulphate  or  aluminates,  which  are  soluble,  are 
formed. 

Aluminum  Hydroxide  A1(OH)3  is  found  in  nature  as  the 
mineral  hydrargillit.  Diaspore  is  a  hydrated  oxide  A12O3-  H2O, 
and  bauxite  is  a  hydrated  oxide  A12O3  •  3  H2O,  containing  ferric 
oxide.  When  an  alkaline  hydroxide  is  added  to  a  solution  of 
an  aluminum  salt,  aluminum  hydroxide  is  precipitated  in  the 
form  of  a  gelatinous  mass :  — 

A1C13  +  3  KOH  =  3  KC1  +  A1(OH)3. 


THE  METALS  OF  THE  EARTHS  461 

Aluminum  hydroxide  is  also  formed  when  an  alkaline  carbonate 
is  added  to  a  solution  of  an  aluminum  salt  :  — 

2  A1C13  +  3  Na2CO3  =  6  NaCl  +  A12(CO3)3,  and 
A12(C03)3  +  3  H20  =  2  A1(OH)3  +  3  CO2. 

The  aluminum  carbonate  formed  is  at  once  completely  decom- 
posed by  hydrolysis,  as  indicated.  Aluminum  hydroxide  may 
be  dehydrated  by  heat.  At  about  100°,  A12O3  -  2  H2O,  and  at 
300°  A12O3  •  H2O  is  formed  ;  finally,  on  strong  ignition  all  the 
water  is  driven  off,  thus  leaving  A12O3. 

The  fact  that  a  carbonate  of  aluminum  does  not  exist  shows 
the  feebly  basic  character  of  aluminum  hydroxide.  Indeed, 
the  latter  is  made  on  a  large  scale  by  fusing  bauxite  with 
soda,  extracting  the  resulting  mass  containing  sodium  alumi- 
nate  with  water  and  passing  carbon  dioxide  into  the  solution, 
thus  :  — 

(1)  Na2CO3  +  A12O3  -  2  H2O  =  2  NaAlO2  +  2  H2O  +  CO2. 

(2)  2  NaA102  +  3  H2O  +  CO2  =  Na2CO3  +  2  A1(OH)3. 

As  the  second  equation  indicates,  sodium  carbonate  and  alu- 
minum hydroxide  are  formed  simultaneously.  The  aluminates 
are  therefore  rather  unstable  salts,  showing  that  aluminum 
hydroxide  is  but  a  feeble  acid.  Just  as  the  zincates  are  formed 
by  dissolving  zinc  hydroxide  in  caustic  alkalies,  so  the  aluminates 
may  be  obtained  by  dissolving  aluminum  hydroxide  in  caustic 
alkalies,  thus  :  — 

Zn(OH)2  +  2  NaOH  =  Na2ZnO2  +  2  H2O. 

A1(OH)3  +  3  NaOH  =  Na3AlO3  +  3  H2O. 

A1(OH)3  +  NaOH  =  NaAlO2  +  2  H2O. 

Like  the  zincates,  the  aluminates  suffer  hydrolysis  to  some 
extent  :  — 

+  NaOH. 


Ammonia,  being  a  weak  base,  does  not  react  with  aluminum 
hydroxide  like  sodium  or  potassium  hydroxide. 

The  aluminates  are  salts  formed  by  replacing  the  hydrogen 
atoms  of  aluminum  hydroxide  by  bases.  When  aluminum 
hydroxide  loses  one  molecule  of  water,  A1O  -  OH  or  HA1O2, 
results  which  may  be  regarded  as  meta-aluminic  acid,  analo- 
gous to  metaboric  acid  BO  •  OH  or  HBO2.  Salts  of  meta< 


462  OUTLINES   OF   CHEMISTRY 

aluminic  acid,  that  is  meta-aluminates,  are  found  in  nature  and 
are  known  as  spinels.  These  minerals  generally  crystallize  in 
regular  octahedra.  Thus  we  have  spinel  Mg(AlO2)2,  galmite 
or  zinc  spinel  Zn(AlO)2,  iron  spinel  or  pleonast  Fe(AlO2)2,  and 
chrysoberyl  G1(A1O2)2.  The  latter  is  rhombic.  The  spinels 
are  all  very  stable.  They  have  been  prepared  artificially  by 
heating  the  constituent  oxides  together,  using  boric  anhydride 
as  a  flux.  Insoluble  aluminates  may  also  be  prepared  by  pre- 
cipitation, so,  for  instance  :  — 

CaCl2  +  2  NaAlO2  =  2  NaCl  +  Ca(AlO2)2. 

It  will  be  recalled  that  calcium  aluminate  is  one  of  the  prod- 
ucts formed  when  cement  sets.  This  aluminate  hardens 
under  water. 

When  aluminum  hydroxide  is  precipitated  in  a  solution  con- 
taining coloring  matter,  the  latter  commonly  unites  with  the 
precipitate,  and  thus  the  supernatant  liquor  remains  clear. 
Suspended  matter  is  also  dragged  down  with  the  precipitate. 
This  fact  is  sometimes  used  in  clarifying  drinking  water.  Pig- 
ments called  lakes  are  made  by  dissolving  dyestuffs  in  water 
together  with  an  aluminum  or  tin  salt,  and  then  adding  an 
alkali  like  soda  ;  thus  precipitates  are  formed  which  consist  of 
the  coloring  matter  united  with  the  hydroxide  of  aluminum  or 
tin. 

Many  dyestuffs  do  not  unite  directly  with  cotton  fibers, 
which,  it  will  be  recalled,  are  practically  cellulose.  Aluminum 
hydroxide  does  unite  with  cotton  fiber,  and  as  dyestuffs  in  turn 
unite  with  aluminum  hydroxide,  the  latter  may  act  as  a  means 
of  fixing  the  dye  to  the  fabric.  This  is  done  by  first  dipping 
the  fabric  in  a  solution  of  an  aluminum  salt,  then  treating  with 
steam,  whereby  aluminum  hydroxide  is  formed  which  adheres 
to  the  fiber,  and  finally  immersing  the  cloth  in  the  dye,  when 
the  latter  unites  with  the  aluminum  hydroxide,  forming  an 
insoluble  compound  which  is  thus  fixed  on  the  goods.  Some- 
times salts  of  tin  are  similarly  employed  in  fastening  dyestuffs 
to  fabrics.  Substances  that  will  serve  this  purpose  are  called 
mordants,  from  the  Latin  word  meaning  to  bite.  Besides  alum 
and  aluminum  acetate,  salts  of  tin,  chromium,  iron,  and  anti- 
mony are  frequently  employed  as  mordants.  Not  all  dyestuffs 
require  the  use  of  mordants,  for  many  unite  with  the  fiber  directly. 


THE  METALS  OF  THE  EARTHS  463 

Again,  though  wool  or  silk  more  frequently  unite  directly  with 
dyestuffs,  they  must,  nevertheless,  be  treated  with"  suitable 
mordants  in  many  cases.  It  must  be  borne  in  mind  that  wool 
and  silk  are  nitrogenous,  organic  compounds  of  animal  origin, 
and  are  quite  different  chemically  from  cotton  and  linen,  which 
consist  mainly  of  cellulose. 

Aluminum  Chloride  A1C13. —  By  dissolving  aluminum  or  its 
hydroxide  in  hydrochloric  acid  and  evaporating  the  solution, 
deliquescent  crystals,  A1C13  •  6  H2O,  form,  which,  on  being 
heated,  give  off  water  and  hydrochloric  acid,  leaving  the  oxide 
behind:  — 

2  A1C18  -  6  H2O  =  A12O3  +  6  HC1  +  3  H2O. 

Anhydrous  aluminum  chloride  is  made  by  heating  aluminum 
in  chlorine  or  by  passing  that  gas  over  a  heated  mixture  of  car- 
bon and  alumina  :  — 

A12O3  4-  3  C  +  3  C12  =  2  A1C18  +  3  CO. 

The  chloride  sublimes  and  thus  may  readily  be  purified.  It  is 
very  hygroscopic,  and  fumes  in  the  air  because  hydrochloric 
acid  is  formed  by  hydrolysis  when  the  salt  is  in  contact  with 
moisture.  The  chloride  is  used  in  the  synthesis  of  organic 
compounds.  With  alkali  chlorides,  it  forms  double  salts  like 
A1C18  •  NaCl  and  A1C13  •  3  KC1.  Aluminum  fluoride  occurs  in 
cryolite,  which  is  an  analogous  double  salt.  Aluminum  bromide 
AlBr3  and  aluminum  iodide  A1I3  are  colorless  salts  which  are 
quite  analogous  to  the  chloride. 

Aluminum  Sulphide  A12S3  is  formed  by  heating  aluminum 
and  sulphur  together.  It  is  a  yellowish  mass  which  is  decom- 
posed by  water,  forming  the  hydroxide  and  hydrogen  sulphide. 
This  behavior  is  similar  to  that  of  magnesium  sulphide.  Be- 
cause of  its  complete  hydrolysis  by  water,  aluminum  sulphide 
is  not  precipitated  when  ammonium  sulphide  is  added  to  a 
solution  of  an  aluminum  salt.  The  precipitate  consists  of 
aluminum  hydroxide. 

Aluminum  Sulphate  A12(SO4)3  is  formed  by  dissolving  the 
hydroxide  in  sulphuric  acid.  On  a  large  scale,  it  is  prepared 
by  treating  kaolin  or  bauxite  with  sulphuric  acid.  On  evapo- 
rating the  solution,  monoclinic  crystals  A12(SO4)3  •  18  H2O 
may  be  obtained.  The  salt  is  readily  soluble  in  water.  On 
heating,  it  loses  both  water  and  sulphur  trioxide,  leaving 


464  OUTLINES   OF  CHEMISTRY 

aluminum  oxide.  Aluminum  sulphate  is  used  as  a  mordant. 
It  is  also  used  in  sizing  paper.  With  the  exception  of  blotting 
paper  and  filter  paper,  all  papers  are  sized  so  that  they  will  be 
smooth  and  not  absorb  ink.  Rosin,  which  consists  essentially 
of  a  complex  organic  acid,  abietic  acid  C44H64O5,  is  dissolved  in 
caustic  soda,  thus  forming  sodium  abietate,  or  resin  soap,  which 
is  added  to  the  paper  pulp,  and  this  mixture  is  then  treated 
with  aluminum  sulphate.  Thus,  sodium  sulphate  and  insoluble 
aluminum  resinate  are  formed.  The  latter  acts  as  a  binder, 
and  gives  the  paper  a  smooth  surface ;  for  under  the  hot  rollers 
the  resinate  is  melted  and  pressed  upon  the  fibers. 

Alums.  —  Aluminum  sulphate  forms  double  salts  with  ammo- 
nium sulphate  and  the  sulphates  of  the  alkali  metals.  These 
double  sulphates  are  readily  obtained  by  adding  a  solution  of 
the  alkali  sulphate  to  one  of  aluminum  sulphate  and  evapo- 
rating. Crystals  are  thus  obtained  which  are  regular  octa- 
hedra.  The  compounds  all  correspond  to  the  general  formula 
M2SO4  -  A12(SO4)3  -  24  H2O,  or  MA1(SO4)2  - 12  H2O,  where 
M  is  either  NH4,  K,  Na,  Cs,  Rb,  Ag,  or  Tl.  These  double  salts 
are  less  soluble  than  aluminum  sulphate.  They  are  called 
alums.  Potassium  alum  K2SO4  •  A12(SO4)3  •  24  H2O  is  common 
alum.  It  is  prepared  on  a  large  scale  by  calcining  the  mineral 
alunite  K2SO4  •  A12(SO4)3  •  4  A1(OH)3  at  about  500°  and  ex- 
tracting the  mass  with  water,  after  having  exposed  the  material 
to  the  action  of  the  air  and  moisture  for  months.  Alum'  is 
also  made  from  clays,  cryolite,  and  bauxite.  Alunite  occurs  in 
large  quantities  in  Hungary  and  in  Italy  near  Rome.  Alum 
dissolves  in  about  8  parts  of  water.  Its  solutions  are  astringent 
and  have  an  acid  reaction.  Solutions  of  aluminum  acetate 
(C2H3O2)3A1  act  similarly.  When  alum  is  heated,  it  melts, 
loses  water,  and  finally  also  some  sulphur  trioxide,  thus  leaving 
a  somewhat  basic  aluminum  potassium  sulphate  behind.  This 
is  commonly  called  burnt  alum.  On  further  heating,  it  is 
converted  into  potassium  sulphate  and  alumina.  Ammonium 
alum  (NH4)2SO4  •  A12(SO4)3  •  24  H2O  is  also  made  on  a  large 
scale.  It  is  cheaper  than  potassium  alum.  Sodium  alum 
Na2S04  •  A12(SO4)3  -  24  H2O  is  more  soluble  than  either  potas- 
sium or  ammonium  alum;  moreover,  it  does  not  crystallize 
readily  and  is  consequently  not  made  commercially.  Alum  is 
used  as  a  mordant,  also  as  an  astringent  in  medicine,  especially 


THE  METALS  OF  THE  EARTHS  465 

as  a  mouth  wash.  The  acetate  is  to  be  preferred  for  the  latter 
purpose. 

In  the  alums,  the  aluminum  may  be  replaced  by  other 
trivalent  elements  like  iron,  chromium,  manganese,  gallium, 
and  indium.  The  compounds  so  obtained  are  isomorphous 
with  the  alums.  They  are  consequently  also  called  alums, 
though  they  contain  no  aluminum.  So  we  have  ferric  ammo- 
nium alum  (NH4)2SO4  .  Fe2(SO4)3  .  24  H2O,  potassium  chrome 
alum  K2SO4  •  Cr2(SO4)3  •  24  H2O,  sodium  manganese  alum 
Na2SO4  •  Mn2(SO4)3  •  24  H2O,  etc.  In  general,  the  formula 

of  an  alum  is  M2SO4  •  M2(SO4)3  •  24H2O,  or  M  .  M(SO4)2  - 

12  H2O,  where  M  is  a  univalent  metal  or  radical,  and  M  is  a 
trivalent  metal. 

Aluminum  Silicates  are  found  in  enormous  quantities  in  nature. 
They  are  also  very  widely  distributed.  Thus  potash  feldspar, 
orthoclase,  K2O  -  A12O3  •  6  SiO2,  or  KAlSi3O8,  soda  feldspar, 
albite,  Na2O  •  A12O3  •  6  SiO2,  or  NaAlSi3O8,  lime  feldspar,  anor- 
thite,  CaO  •  A12O3  •  2  SiO2,  or  CaAl2Si2O8,  occur  in  granitic 
rocks,  together  with  quartz  and  micas.  The  latter  are  also 
silicates  of  aluminum,  containing  potassium,  magnesium,  cal- 
cium, and  sometimes  iron.  The  mineral  disthen  is  a  pure 
aluminum  silicate  Al2SiO5,  which  occurs  in  rhombic  crystals; 
it  is  also  found  in  triclinic  forms  as  andalusite.  These  two 
minerals  are  rather  rare. 

By  the  action  of  water  and  carbon  dioxide  of  the  air  upon  feld- 
spars, the  latter  lose  their  alkali  content,  which  is  dissolved  and 
enters  the  soil  as  soluble  silicates  or  carbonates,  thus  supplying 
potash  and  lime  needed  for  plant  growth.  A  hydrous  aluminum 
silicate  H2  Al2(SiO4)2  •  2  H2O,  called  kaolin,  remains  as  a  white 
clay.  The  reactions  of  the  weathering  process  of  typical  feld- 
spars are  in  the  main  as  follows  :  — 

2  KAlSi308  +  2  H20  =  K2Si409  +  H2Al2(SiO4)2  -  H2O. 
Ca  Al2Si2O8  +  2  CO2  +  3  H2O  =  Ca(HCO3)2  +  H2  Al2(SiO4)2  -H2O. 

Ordinary  clays  contain  ferric  hydroxide,  sand,  and  various  sili- 
cates besides  kaolin.  Frequently  calcium  carbonate  is  also 
present.  Clay  containing  a  large  amount  of  calcium  carbonate 
is  called  marl. 

When  mixed  with  water,  clay  forms  a  plastic  mass  that  can 
readily  be  molded  into  various  forms  as  desired.  On  drying 


466  OUTLINES  OF  CHEMISTRY 

and  heating  this  material  to  higher  temperatures  (i.e.  "burn- 
ing "  or  "  firing  "  it),  it  becomes  hard,  dense,  and  resistant  with- 
out actually  melting.  These  facts  form  the  basis  of  making 
bricks,  earthenware,  and  porcelain  from  clay.  The  rather  im- 
pure clays  are  used  for  the  manufacture  of  bricks,  which  get 
their  color  from  the  iron  oxides  contained  in  the  clay.  Red 
pottery  is  made  from  similar  material.  Light-colored  stone- 
ware and  pottery  requires  a  clay  relatively  free  from  iron; 
while  fine  white  porcelain  necessitates  kaolin  that  contains  no 
iron.  Earthenware  is  frequently  porous.  Sometimes  this 
porous  material  is  covered  with  a  glaze ;  and  again,  as  porce- 
lain, the  material  may  be  made  perfectly  vitreous  throughout. 
Porous  earthenware  is  made  by  simply  shaping  clay  containing 
but  little  fusible  material  and  firing  it,  as  in  making  bricks, 
flowerpots,  etc.  If  such  articles  are  to  be  glazed,  they  are 
covered  with  a  layer  of  fusible  silicates,  which  is  fired  on.  In 
the  case  of  the  cheaper  grades  of  stoneware,  glazed  bricks, 
etc.,  this  is  accomplished  by  simply  throwing  salt  into  the  kiln. 
At  the  high  temperature  that  obtains,  the  salt  is  volatilized  and 
decomposed  as  it  comes  in  contact  with  the  earthenware ; 
hydrochloric  acid  and  a  readily  fusible  sodium  aluminum  sili- 
cate being  formed.  The  latter  covers  the  surface  of  the  arti- 
cles as  a  glaze.  In  making  porcelain,  the  purest  kaolin  is  mixed 
with  finely  pulverized  feldspar  and  quartz.  Of  this  mixture 
the  dishes  are  shaped,  and  when  they  are  finally  fired  in  the 
kiln,  the  feldspar  melts,  fills  the  pores,  and  thus  produces  the 
translucent  material  known  as  porcelain.  In  actual  practice, 
the  articles  are  fired  twice.  After  the  first  heating,  they  are 
dipped  in  water  containing  in  suspension  very  finely  ground 
material  intended  for  the  glaze.  This  material  commonly  con- 
sists of  kaolin,  together  with  enough  finely  ground  feldspar  to 
make  a  mixture  that  will  fuse  at  a  somewhat  lower  temperature 
than  the  ware  of  which  the  articles  are  composed.  This  fine 
material  is  taken  up  by  the  porous  porcelain,  which  is  then 
dried  and  fired  again.  As  the  feldspar  melts,  a  smooth  glaze 
is  produced,  and  a  mass  that  is  vitreous  throughout  results. 

Ultramarine  is  an  important  blue  pigment  which  is  manu- 
factured in  large  quantities  by  heating  together  soda,  sulphur, 
and  clay,  or,  more  frequently,  sodium  sulphate,  carbon,  and 
clay.  The  product,  which  is  a  silicate  of  sodium  and  aluminum 


THE  METALS  OF  THE  EARTHS  467 

combined  with  sodium  sulphides,  is  at  first  green  and  is  called 
ultramarine  green.  This  is  also  used  as  a  pigment.  On  heat- 
ing ultramarine  green  in  a  current  of  air,  it  turns  blue,  forming 
ultramarine  blue.  Ultramarine  violet  results  when  the  blue  is 
heated  in  a  current  of  hydrochloric  acid  gas  to  about  175°,  and 
ultramarine  red  is  formed  by  the  same  process  at  about  145°. 
Only  the  green  and  blue  ultramarines  have  much  value  in 
practice.  The  constitution  of  ultramarine  has  caused  much 
discussion,  and  has  not  yet  been  finally  settled.  It  is  probable 
that  ultramarine  blue  is  (Na2Al2Si2O8)2  •  Na2S2.  It  is  found  in 
nature  as  lapis  lazuli,  from  which  the  pigment  was  first  prepared. 
Ultramarine  is  stable  towards  light,  and  the  action  of  the  air 
and  water ;  but  acids  readily  destroy  the  color,  liberating  sul- 
phur and  hydrogen  sulphide.  Ultramarine  blue  is  much  used 
in  paints  and  laundry  blue.  It  is  also  employed  in  removing 
the  yellow  tinge  of  sugar,  paper  pulp,  linen  and  cotton  fabrics, 
and  in  making  cotton  prints  and  wall  paper. 

Analytical  Tests  for  Aluminum.  —  Aluminum  is  precipitated 
as  the  gelatinous  hydroxide  from  solutions  by  either  potassium, 
sodium,  or  ammonium  hydroxides,  carbonates,  or  sulphides. 
Aluminum  hydroxide  is  soluble  in  potassium  or  sodium  hydrox- 
ides, but  only  slightly  soluble  in  ammonium  hydroxide.  By 
means  of  hydrogen  sulphide  or  carbon  dioxide,  solutions  of 
aluminates  are  decomposed,  yielding  a  precipitate  of  aluminum 
hydroxide.  When  alumina  is  moistened  with  a  solution  of 
cobalt  nitrate  and  then  strongly  ignited,  a  blue  mass 
Co(AlO2)2  results.  This  blue  is  known  as  Thenard's  blue  or 
cobalt  ultramarine. 

Gallium.  — This  is  a  rare  metal  whose  atomic  weight  is  69.9. 
It  was  discovered  in  1875  by  Lecoq  de  Boisbaudran,  who 
detected  its  presence  in  a  sample  of  zinc  blende  from  Pierrefitte 
in  the  Pyrenees,  by  means  of  the  spectroscope.  In  1869  Men- 
deleeff  foretold  the  existence  and  described  the  properties  of 
this  element,  which  he  called  eka-aluminum.  He  was  able  to 
do  this  from  the  periodic  system  of  the  elements,  in  which  the 
space  now  occupied  by  gallium  was  then  vacant.  Gallium 
receives  its  name  from  Gaul,  the  native  land  of  its  discoverer. 
Gallium  is  a  hard,  white  metal  which  is  stable  in  the  air.  It 
melts  at  30°,  and  its  specific  gravity  is  5.9.  In  its  compounds, 
it  is  trivalent  and  analogous  to  aluminum,  though  a  chloride  of 


468  OUTLINES   OF  CHEMISTRY 

the  formula  GaCl2  is  also  known.  The  spark  spectrum  of  gal- 
lium exhibits  two  characteristic  lines  in  the  violet,  through 
which  the  element  was  discovered. 

Indium,  whose  atomic  weight  is  114.8,  is  also  a  rare  metal. 
It  was  discovered  by  Reich  and  Richter  in  1863  by  means  of  the 
spectroscope,  in  a  sample  of  zinc  blende  from  Freiberg.  The 
metal  receives  its  name  from  the  bright  indigo-blue  line  that 
characterizes  its  spectrum.  Indium  is  a  silver-white  metal, 
softer  than  lead  and  very  malleable.  Its  specific  gravity  is  7.4, 
and  its  melting  point  is  176°.  Like  gallium,  it  is  stable  in  the 
air.  In  its  compounds,  it  is  commonly  trivalent  like  aluminum, 
though  indium  dichloride  InCl2  and  monochloride  InCl  are 
also  known. 

Thallium  and  its  Compounds.  —  Thallium  is  a  metal  whose 
physical  properties  are  similar  to  those  of  lead.  It  was  dis- 
covered in  1861  by  means  of  the  spectroscope  by  Sir  William 
Crookes,  in  the  mud  at  the  bottom  of  the  lead  chambers  of  the 
sulphuric  acid  factory  at  Tilkerode  in  the  Harz.  Thallium 
compounds  yield  a  very  characteristic  green  line  in  the  spec- 
trum, whence  the  name  thallium,  meaning  a  green  branch. 
Crookesite,  a  selenide  of  copper,  thallium,  and  silver  contains 
about  17  per  cent  thallium.  Many  native  sulphides  of  iron  and 
copper  contain  small  amounts  of  thallium,  whence  its  appear- 
ance in  the  flues  and  lead  chambers  of  sulphuric  acid  plants. 
In  1862  Lamy  demonstrated  that  thallium  is  a  metal.  It  oxi- 
dizes readily  on  exposure  to  air,  and  hence  is  kept  under  petro- 
leum oil  or  glycerine.  The  atomic  weight  of  thallium  is  204.0, 
and  its  valence  is  either  1  or  3. 

In  thallous  compounds,  the  metal  is  univalent  and  therefore 
analogous  to  the  alkali  metals.  Thus  we  have  T1OH,  T12O, 
T1F,  T1C1,  TIBr,  Til,  T1C1O3,  T1NO3,  T12CO3,  T12SO4,  T12S, 
Tl2PtCl6,  etc.  Thallous  chloride,  bromide,  and  iodide  are 
nearly  insoluble  in  water,  and  so  in  this  respect  these  com- 
pounds are  similar  to  the  corresponding  ones  of  silver.  Thai- 
lie  chloride  T1C13  is  very  soluble  in  water. 

In  thallic  compounds,  thallium  is  trivalent.  Thus  we  have 
T10  •  OH,  T1203,  T1C13,  T1(N03)3,  T12(SO4)3,  T12S3,  etc. 

Thallium  readily  dissolves  in  nitric  or  sulphuric  acid,  whereas, 
on  account  of  the  fact  that  thallous  chloride  is  difficultly 
soluble,  hydrochloric  acid  attacks  the  metal  but  slightly,  it 


THE  METALS  OF  THE  EARTHS  469 

being  soon  covered  with  a  protective  coating  of  the  chloride. 
Thallium  salts  are  readily  recognized  by  means  of  the  spectro- 
scope. From  neutral  or  faintly  acid  solutions  of  thallous  salts, 
hydrogen  sulphide  precipitates  black  thallous  sulphide,  which 
is  insoluble  in  water  and  acetic  acid,  but  it  readily  dissolves  in 
sulphuric  or  nitric  acid. 

The  Rare-Earth  Elements.  — These  are  a  series  of  metals  that 
form  compounds  which  are  in  general  analogous  to  those  of 
aluminum.  The  rare-earth  metals  are,  as  the  name  implies,  of 
very  rare  occurrence.  They  are  found  in  the  complex  and 
very  rare  minerals  monazite,  gadolinite,  euxenite,  samarskite, 
orthite,  cerite,  yttrotantalite,  hjelmite,  and  several  others  which 
occur  mainly  in  the  Scandinavian  peninsula,  Finland, -Green- 
land, France,  Bavaria,  and  the  United  States.  In  general,  the 
elements  are  trivalent,  like  aluminum. 

Chemically  the  rare-earth  metals  deport  themselves  very 
much  alike,  which  makes  it  extremely  difficult  to  separate 
them.  In  general,  their  oxalates  are  insoluble,  and  their  sul- 
phates are  soluble  and  readily  form  double  salts  with  the  sul- 
phates of  the  alkalies.  As  the  nitrates  are  decomposed  by  heat 
at  different  temperatures,  this  fact  is  used  in  separating  the 
rare  earths  from  one  another.  Fractional  crystallization  and 
fractional  precipitation  are  also  used  to  effect  separations  ;  but 
these  processes  are  laborious,  and  generally  yield  products  that 
are,  after  all,  not  quite  pure. 

Scandium  (Sc — 44.1)  was  discovered  in  1879  by  Nilson  and 
Cleve.  Mendeleeff  predicted  the  existence  of  this  element  in 
1869.  Compare  gallium  and  germanium.  He  called  it 
ekaboron,  and  described  its  properties.  The  element  occurs  in 
euxenite  and  gadolinite.  Volatilized  in  the  electric  arc,  scan- 
dium chloride  yields  a  bright  characteristic  spectrum  of  many 
lines.  The  oxide  Sc2O3  is  a  white,  earthy  powder. 

Yttrium  (Y — 89.0)  occurs  in  the  silicate  gadolinite  found  at 
Ytterby,  whence  its  name.  Yttrium  was  discovered  in  1843 
by  Mosander.  Its  oxide  Y2O3  is  a  white,  earthy  powder  of 
very  high  melting  point.  The  chloride  YC13  yields  a  bright 
spectrum  containing  many  lines. 

Lanthanum  (La  — 139.0)  was  found  in  cerite  in  1839  by 
Mosander.  The  name  lanthanum  comes  from  the  Greek,  mean- 
ing hidden.  Like  yttrium,  lanthanum  may  be  prepared  by 


470  OUTLINES   OF   CHEMISTRY 

electrolysis  of  its  molten  chloride.  Lanthanum  is  a  white, 
malleable  metal  not  unlike  iron  in  appearance.  When  heated, 
it  burns  to  oxide  La2O5,  a  white  powder  which  readily  absorbs 
water,  forming  La(OH)3.  Other  compounds  are  La2(CO3)3  • 
3H2O,  La(NO3)3  -  6  H2O,  La2(SO4)3  •  9  H2O,  LaCl3.  The 
latter  is  deliquescent  and  readily  forms  2  LaCl3  •  15  H2O  ;  it 
shows  a  characteristic  spectrum  of  many  lines. 

Ytterbium  (neoytterbium)  (Yb  — 172.0)  occurs  with  scandium 
and  yttrium  in  euxenite  and  gadolinite.  Its  compounds  are 
in  general  like  those  of  yttrium.  In  1907  Urbain  in  Paris 
showed  that  the  old  ytterbium  consists  of  two  elements,  neoyt- 
terbium (Yb^-172.0)  and  lutecium  (Lu  — 174,0).  Auer  von 
Welsbach,  in  Vienna,  published  the  same  discovery  almost  simul- 
taneously. He  named  the  two  new  elements  aldebaranium  and 
cassiopeium,  and  found  their  atomic  weights  to  be  172.90  and 
174.23,  respectively.  As  Urbain's  work  appeared  first,  his 
nomenclature  will  probably  be  adopted. 

Cerium  (Ce — 140.25)  resembles  lanthanum  very  much.  It 
was  discovered  by  Klaproth  and  also  by  Berzelius  and  Hisinger 
in  1803.  The  element  is  named  from  the  planet  Ceres,  which 
had  just  been  discovered.  .  The  mineral  cerite  contains  about 
60  per  cent  cerium.  Cerium  forms  compounds  in  which  the 
element  is  trivalent,  like  Ce2O3,  Ce2(SO4)3,  and  CeCl8,  but 
it  also  forms  compounds  in  which  it  is  quadrivalent,  like  CeO2, 
Ce2(SO4)2  -  4  H2O,  Ce(OH)4,  CeF4  .  H2O.  The  latter  class  of 
compounds  indicates  that  cerium  is  closely  related  to  other 
quadrivalent  elements,  and  consequently  probably  belongs  in 
the  same  group  as  silicon,  titanium,  zirconium,  and  thorium. 
Cerium  may  be  prepared  by  electrolysis  of  the  molten  chloride. 
It  is  a  steel-gray,  very  malleable  metal  of  specific  gravity  7.0. 
It  readily  burns  in  the  air,  forming  a  yellow  powder,  CeO2. 
Ceric  hydroxide  Ce(OH)4  is  a  red  precipitate.  In  general, 
cerous  compounds  are  colorless,  while  the  eerie  compounds  are 
yellow,  brown,  or  red.  Cerium  nitrate  is  now  prepared  from 
monazite  sand  found  in  North  Carolina  in  considerable  quan- 
tities. It  is  used  together  with  thorium  nitrate  in  making  the 
mantles  for  Welsbach  gas  lights.  The  process  of  making  these 
mantles  consists  essentially  of  knitting  the  mantle  of  cotton 
thread,  and  then  dipping  this  into  solutions  of  the  nitrates  of 
cerium  and  thorium.  The  mantle  is  then  dried  and  calcined; 


THE  METALS  OF  THE  EARTHS  471 

thus  the  thread  is  destroyed,  and  the  nitrates  are  converted  to 
oxides  which  adhere  together.  Such  mantles  are,  of  course, 
always  frail,  and  consequently  must  be  handled  with  care.  As 
already  mentioned,  experience  has  shown  that  mantles  consist- 
ing of  1  per  cent  cerium  oxide  CeO2  and  99  per  cent  thorium 
oxide  ThO2  give  the  best  efficiency.  The  light  emitted  from 
such  mantles,  when  heated  to  incandescence  by  a  Bunsen 
burner,  is  a  brilliant  white.  Mantles  containing  different  pro- 
portions of  ceria  and  thoria  are  less  efficient,  and  those  made  of 
other  earths  give  a  light  of  poorer  quality  and  also  lower  inten- 
sity. The  Nernst  lamp  consists  of  a  filament  of  earths  heated 
to  incandescence  by  passing  an  electric  current  through  it.  At 
ordinary  temperatures,  this  filament  is  practically  a  non-con- 
ductor ;  but  on  being  heated,  it  conducts  electricity  and  gives 
,  a  bright  light. 

Praseodymium  (Pr— 140.6)  and  neodymium  (Nd— 144.3) 
were  for  half  a  century  regarded  as  but  one  element,  didymium. 
But  in  1885  Auer  von  Welsbach  separated  it  into  praseodym- 
ium, which  forms  leek-green  salts,  and  neodymium,  which  forms 
rose-violet  salts.  These  elements  occur  with  cerium  and  lan- 
thanum, which  they  resemble.  The  absorption  spectra  of  solu- 
tions of  salts  of  the  didymiums  are  especially  characteristic. 

Samarium  (Sa — 150.4)  was  discovered  in  Samarskite,  .  a 
mineral  from  North  Carolina,  in  1878  by  Delafontaine  and  also 
by  Lecoq  de  Boisbaudran.  Typical  compounds  are  Sm2O3, 
SmCl3.6H20,  Sm2(S04)3  .-8  H2O,  Sm(NO3)2  .  6  H2O.  These 
compounds  are  similar  to  those  of  lanthanum. 

Erbium  (Er— 167.7),  terbium  (Tb— 159.2),  thulium  (Tm— 
168.5),  and  dysprosium  (Dy — 162.5)  are  found  associated  with 
yttrium.  Terbium  also  occurs  in  samarskite  and  in  small  quan- 
tities in  gadolinite,  which  mineral  also  contains  gadolinium 
(Gd  — 157.3).  Europium  (Eu  — 152.0)  is  another  element 
about  which  but  little  is  known. 

The  position  of  a  goodly  number  of  these  elements  in  the 
periodic  system  is  still  unsettled.  It  is  also  quite  probable 
that  some  of  the  rare-earth  elements,  especially  those  which  have 
been  studied  but  little,  will  be  separated  further. 


CHAPTER  XXVII 

LEAD   AND   TIN 

GERMANIUM,  lead,  and  tin  are  three  metallic  elements  that 
exhibit  a  valence  of  two  or  four  in  their  compounds,  which  are 
consequently  analogous  to  those  of  carbon  and  silicon.  The 
hydroxides  of  these  metals  increase  in  basicity  with  rise  in 
atomic  weight.  These  hydroxides,  however,  also  have  weakly 
acidic  properties,  for  towards  strong  bases  they  may  act  as 
acids.  Germanium  forms  a  compound  with  hydrogen,  GeH4, 
but  tin  and  lead  do  not. 

Germanium  (Ge  —  72.5)  is  a  very  rare  metal  whose  existence 
was  foretold  by  Mendeleeff  in  1871.  He  called  the  element 
ekasilicon  and  described  its  properties  and  fixed  its  atomic 
weight  at  about  73.  In  1886  Clemens  Winkler  discovered 
germanium  in  the  silver  ore  argyrodite  GeS2  •  3  Ag2S,  which 
occurs  near  Freiberg.  The  metal  is  also  found  in  euxenite, 
samarskite,  and  confieldite.  Germanium  has  properties  which 
are  practically  the  same  as  those  predicted  by  Mendeleeff  for 
ekasilicon.  Germanium  is  a  grayish  white,  brittle  metal  of 
specific  gravity  5.47.  It  crystallizes  in  octahedra,  and  melts 
at  about  900°.  From  the  formulae  of  the  following  compounds, 
the  analogy  of  germanium  to  carbon  and  silicon  is  evident : 
GeH4,  GeF4,  GeCl4,  GeHClg,  GeO2,  GeOCl2,  GeS2,  K2GeF6. 
Besides  these,  Ge(OH)2  and  GeS  are  also  known. 

Occurrence,  Metallurgy,  and  Properties  of  Tin.  —  Tin  occurs 
mainly  in  cassiterite,  or  tin  stone,  SnO2,  which  crystallizes  in 
the  tetragonal  s}^stem  (Fig.  59),  and  is  usually  colored  brown 
or  black  by  oxides  of  iron  and  manganese.  The  metal  has 
been  known  for  a  long  time,  for  though  it  does  not  occur  in 
the  free  state,  its  oxide  is  readily  reduced  by  means  of  carbon. 
Tin  was  alloyed  with  copper  to  make  bronze,  even  in  ancient 
times.  The  Latin  name  of  tin  is  stannum,  whence  the  symbol 
Sn.  It  was  also  called  plumbum  candidum  to  distinguish  it 
from  lead,  plumbum  nigrum.  Tin  was  obtained  from  the 

472 


LEAD  AND   TIN  473 

mines  at  Cornwall  in  the  days  of  the  Roman  Empire.  Tin 
ores  also  occur  in  Saxony,  Peru,  Australia,  Alaska,  and  the 
islands  of  Billiton  and  Banca  east  of  Sumatra. 

To  extract  tin  from  its  ores,  the  latter  are  first  roasted  to 
expel  any  sulphur  or  arsenic  present.  The  material  is  then 
treated  with  crude  hydrochloric  acid  to  remove  iron,  copper, 
etc.,  and  lixiviated  with  water,  after  which  the  finely  divided 
ore  is  mixed  with  carbon  and  heated  in  a  furnace,  thus  :  — 


The  molten  tin  collects  at  the  bottom  of  the  furnace  and  is 
drawn  off  and  cast  into  bars.  It  is  purified  by  remelting  it 
and  collecting  that  portion  which  fuses  at  the  lowest  tempera- 
ture, for  this  is  freest  from  other  metals.  Banca  tin  is  the 
purest  ;  though  German  and  English  tin,  block  tin,  also  often 
is  98  to  99.9  per  cent  pure.  The  world's  annual  output  of  tin 
is  about  110,000  tons,  most  of  which  comes  from  the  East 
Indies. 

Tin  is  a  silver-  white,  lustrous  metal  of  specific  gravity  7.3. 
It  melts  at  232°,  and  boils  at  about  1600°.  Tin  is  crystalline 
in  character,  and  its  bars,  when  bent,  give  a  low,  peculiar,  crac- 
kling noise  known  as  tin  cry^  which  is  due  to  the  friction  of 
the  crystalline  particles  moving  over  one  another.  Tin  is  very 
ductile,  malleable,  and  so  soft  that  it  may  readily  be  cut  with 
a  knife.  At  100°,  the  metal  is  still  malleable,  but  this  property 
decreases  on  raising  the  temperature  further.  At  about  200°, 
tin  is  brittle  and  may  be  powdered.  On  cooling,  molten  tin 
always  congeals  in  crystalline  form.  At  low  temperatures,  tin 
gradually  changes  to  a  gray,  brittle  variety.  This  change 
takes  place  most  rapidly  at  —  48°,  though  it  proceeds  appre- 
ciably even  at  —15°.  It  is  the  cause  of  the  tin  pest,  which 
consists  of  the  disintegration  of  tin  organ  pipes  and  tin  roofs 
and  gutters  in  Russia,  where  winters  are  very  cold.  Above  20°, 
the  ordinary  malleable  tin  is  the  stable  form,  while  below  that 
temperature  the  gray,  brittle  variety  is  the  stable  modification. 

In  the  air  tin  remains  practically  unchanged.  When 
strongly  heated,  it  burns  with  a  white  flame  to  SnO2.  In  hot 
hydrochloric  acid  tin  dissolves  :  — 


Sn+2HCl=SnC 


474  OUTLINES  OF  CHEMISTRY 

Hot  concentrated  sulphuric  acid  acts  on  tin,  forming  sulphut 
dioxide  and  stannous  sulphate,  SnSO4:  — 

Sri  +  2  H2SO4  =  SnSO4  +  2  H2O  +  SO2. 

Cold,  dilute,  nitric  acid  acts  on  tin,  forming  stannous  nitrate, 
thus:  — 

4  Sn  + 10  HN03  =  4  Sn(NO3)2  +  3  H2O  +  NH4NO3. 

Concentrated  nitric  acid  converts  tin  into  metastannic  acid 
(which  see). 

Uses  of  Tin. — Tin  is  used  to  a  very  large  extent  in  making 
tin  plate  or  tinned  iron.  This  process  consists  of  dipping  thor- 
oughly cleaned  sheet  iron  in  molten  tin.  The  ordinary  tin 
cans  and  other  tin  utensils  are  made  of  iron  covered  with  a  layer 
of  tin  by  this  dipping  process.  Copper  is  also  often  tinned  in 
the  same  way. 

Solder  usually  consists  of  1  part  tin  and  1  part  lead,  but 
these  proportions  are  often  varied.  An  alloy  of  about  90  per 
cent  tin,  8  per  cent  antimony,  and  2  per  cent  copper  is  called 
Britannia  metal.  Pewter  consists  of  75  per  cent  tin  and  25  per 
cent  lead.  Bronzes  contain  tin,  copper,  and  sometimes  also 
zinc,  as  stated  under  copper.  Alloys  of  silver  and  tin  are 
employed  to  make  amalgams  for  filling  teeth.  Tin  is  recov- 
ered from  old  tin  cans,  etc.,  either  by  melting  off  the  coating 
or  more  frequently  by  electrolysis.  In  this  process,  the  cans 
placed  in  a  wire  basket  are  the  anode,  caustic  alkali  solution 
serves  as  the  electrolyte,  and  an  iron  plate  is  used  as  the 
cathode.  Much  tin  is  saved  in  this  way. 

Chlorides  of  Tin.  —  Stannous  chloride  SnCl2  is  formed  by  dis- 
solving the  metal  in  hydrochloric  acid.  From  the  solution, 
monoclinic  crystals  SnCl2-  2  H2O  are  obtained,  which  are  very 
soluble  in  water.  This  solution  is  a  strong  reducing  agent. 
For  instance,  it  readily  reduces  mercuric  chloride  to  calomel 
and  even  to  metallic  mercury:  — 

2  HgCl2  +  SnCl2  =  2  HgCl  +  SnCl4. 
HgCl2  +  SnCl2  =  Hg  +  SnCl4. 

Stannous  chloride  is  hydro!  yzed  in  solution.  By  treatment 
with  much  water  a  basic  chloride  of  the  composition  Sn(OH)Cl 
is  precipitated :  — 

SnCl2  +  H2O  =  Sn(OH)Cl  +  HC1. 


LEAD   AND   TIN  475 

Stannous  chloride  is  used  as  a  reducing  agent  in  the  laboratory 
and  as  a  mordant  in  dyeing  fabrics. 

Stannic  chloride  SnCl4  is  made  by  the  action  of  chlorine  on 
tin  or  stannous  chloride,  or  by  treating  stannic  oxide  or 
hydroxide  with  hydrochloric  acid.  It  is  a  colorless,  fuming 
liquid  of  specific  gravity  2.28.  It  boils  at  114°,  and  congeals 
at  —  33°.  Tin  tetrachloride  is  also  known  as  spiritus  fumans 
Libavii.  With  water,  the  chloride  forms  the  hydrates 
SnCl4-3HaO,  butter  of  tin,  SnCl4.5H2O,  and  SnCl4 . 8  H2O. 
These  are  soluble  in  water.  Stannic  chloride  readily  forms 
double  salts  with  other  chlorides,  like  SnCl4  •  2  HC1  •  6  H2O,  or 
H2SnCl6  •  6  H2O ;  SnCl4  •  2  KC1,  or  K2SnCl6 ;  SnCl4  •  2  NH4C1,  or 
(NH4)2SnCl6.  The  latter  is  called  pink  salt.  It  is  used  as 
a  mordant,  as  is  also  the  hydrate  SnCl4  •  5  H2O.  Stannic 
chloride^,  on  being  boiled  with  water,  yields  a  precipitate  of 
stannic  acid,  thus  : — 

SnCl4  +  3  H2O  =  4  HC1  +  H2SnO3. 

In  hydrocarbons,  carbon  disulphide,  and  many  other  organic 
and  inorganic  liquids  SnCl4  is  soluble  in  all  proportions. 
Compounds  like  SnCl4-PCl6,  SnCl4.2SCl4,  SnCl4-2NOCl, 
SnCl4  •  POC13  are  also  known.  Indeed,  stannic  chloride  enters 
into  a  large  number  of  compounds. 

Tin  tetrabromide  SnBr4  and  tin  tetraiodide  SnI4  are  also 
known. 

Oxides  of  Tin.  —  Stannous  oxide  SnO  is  formed  as  a  black 
powder  by  heating  stannous  hydroxide  out  of  contact  with 
oxygen.  When  ignited  in  the  air,  it  burns,  forming  stannic 
oxide  SnO2.  Stannous  hydroxide  is  formed  by  adding  sodium 
carbonate  to  a  stannous  chloride  solution :  — 

SnCla  +  Na2CO3  +  H2O  =  2  NaCl  +  Sn(OH)2  +  CO2. 

In  potassium  or  sodium  hydroxide,  stannous  hydroxide  is  sol- 
uble, but  not  in  ammonium  hydroxide,  thus  :  — 

Sn(OH)2  +  2  KOH  =  K2SnO2  +  2  H2O. 

On  boiling  the  solution,  the  potassium  or  sodium  stannite  is 
converted  to  stannate,  with  concomitant  precipitation  of  tin, 
thus : — 

2  K2Sn02  +  H20  =  K2SnO3  +  2  KOH  +  Sn. 


476  OUTLINES   OF   CHEMISTRY 

Stannic  oxide  SnO2,  or  stannic  acid  anhydride,  is  obtained  by 
burning  tin  in  oxygen  or  in  the  air,  or  by  igniting  stannic  acid. 
Stannic  oxide  is  found  in  nature  in  quadratic  crystals  as  cassit- 
erite,  the  principal  ore  of  tin.  Stannic  oxide  as  artificially  pre- 
pared is  a  white  or  slightly  yellowish  powder,  which  is  insoluble 
in  water,  acids,  or  alkalies  especially  after  strong  ignition.  Dur- 
ing ignition,  the  oxide  turns  darker  in  color,  but  it  again  changes 
to  white  on  cooling.  When  fused  with  caustic  alkalies,  stannic 
oxide  forms  stannates,  which  are  soluble,  thus  :  — 

SnO2  +  2  NaOH  =  Na2SnO3  -f  H9O. 

Stannic  hydroxide  is  formed  as  a  gelatinous  precipitate  by 
boiling  tin  tetrachloride  with  water,  or  by  adding  ammonia  to 
the  solution.  Stannic  hydroxide  or  orthostannic  acid  Sn(OH)4 
easily  splits  off  water,  and  forms  SnO  •  (OH)2  or  H2SnQg,  which 
is  called  stannic  acid.  This  also  gradually  loses  water  so  that 
neither  Sn(OH)4  nor  SnO  •  (OH)2  have  really  been  definitely 
isolated.  It  will  be  observed  that  stannic  acid  H2SnO3  is  anal- 
ogous to  H2CO3  and  H2SiO3.  Stannic  acid  readily  dissolves 
in  concentrated  sulphuric,  nitric,  or  hydrochloric  acid,  and  also 
in  dilute  solutions  of  caustic  alkalies.  With  the  latter  it 
forms  stannates,  which  dissolve.  From  these  soluble  stannates, 
the  stannates  of  other  metals  may  be  prepared  by  precipitation  ; 
for  as  in  the  case  of  the  silicates  and  carbonates,  only  the  stan- 
nates of  the  alkalies  are  soluble.  Sodium  stannate  Na2SnO3  is 
made  by  fusing  cassiterite  with  caustic  soda  or  by  fusing  tin 
with  sodium  carbonate  and  sodium  nitrate.  From  its  aqueous 
solutions,  sodium  stannate  separates  in  monoclinic  crystals 
Na2SnO3  •  3  H2O.  It  is  used  as  a  mordant  in  calico  printing, 
being  termed  preparing  salt. 

Concentrated  nitric  acid  converts  tin  into  a  white,  insoluble 
powder  known  as  metastannic  acid,  which  is  a  hydrated  oxide  of 
tin,  whose  composition,  like  that  of  stannic  acid,  varies  slightly 
according  to  temperature  and  other  conditions  of  preparation, 
being  probably  H2SnO4  or  H2SnO3,  thus  :  — 

Sn  +  4  HNO3  =  H2SnO3  +  H2O  +  4  NO2. 

But  metastannic  acid  is  different  from  stannic  acid,  for  it  is 
insoluble  in  acids,  and  with  alkalies  it  forms  salts  like 
K2Sn6On  •  4  H2O  and  Na2Sn5On  •  4  H2O,  metastannates,  which 


LEAD   AND   TIN  477 

would  indicate  that  the  acid  is  dibasic  and  probably  analogous 
to  polysilicic  acids.  On  fusion  with  caustic  alkalies,  metastannic 
acid  yields  the  stannates  which  are  identical  with  those  of  stannic 
acid. 

Sulphides  of  Tin.  —  Stannous  sulphide  SnS  is  formed  as  a 
dark  brown  precipitate  when  hydrogen  sulphide  is  conducted 
into  a  solution  of  a  stannous  salt.  This  precipitate  is  insoluble 
in  dilute  acids  or  solutions  of  the  monosulphides  of  the  alkalies  ; 
but  when  boiled  with  the  latter  and  sulphur,  or  when  treated 
with  alkaline  polysulphides,  soluble  sulpho-stannates  are  formed, 

thus : — 

SnS  +  S  +  K2S  =  K2SnS3. 

SnS  +  (NH4)2S2  =  (NH4)2SnS8. 

These  sulpho-salts  are  decomposed  by  hydrochloric  acid,  yield- 
ing precipitates  of  stannic  sulphide  SnS2,  thus  :  — 

K2SnS8  +  2  HC1  =  2  KC1  +  SnS2  +  H2S. 
(NH4)2SnS3  +  2  HC1  =  2  NH4C1  +  SnS2  +  H2S. 

Stannic  sulphide  is  a  yellow,  amorphous  powder,  which  is  also 
obtained  by  passing  hydrogen  sulphide  into  solutions  of  stannic 
salts.  It  is  insoluble  in  dilute  acids.  On  heating,  it  decom- 
poses into  stannous  sulphide  and  sulphur.  In  alkaline  sulphides 
it  dissolves,  yielding  sulphostannates.  Concentrated  hydro- 
chloric acid  dissolves  stannic  sulphide,  and  concentrated  nitric 
acid  converts  it  into  metastannic  acid. 

By  heating  tin  and  sulphur  together,  stannous  sulphide  is 
formed.  Stannic  sulphide  cannot  be  thus  obtained,  for  it  de- 
composes into  sulphur  and  stannous  sulphide  at  the  high  tem- 
perature reached  during  the  progress  of  the  reaction.  However, 
by  heating  together  sulphur,  ammonium  chloride,  and  finely 
divided  tin,  stannic  sulphide  is  obtained  as  a  golden  yellow, 
crystalline  mass  which  is  used  for  bronzing,  being  called  mosaic 
gold. 

Analytical  Tests  for  Tin.  —  When  mixed  with  sodium  car- 
bonate and  ignited  on  charcoal,  tin  compounds  are  reduced, 
yielding  bright  globules  of  the  metal.  The  behavior  of  stan- 
nous chloride  toward  mercuric  chloride  is  often  used  to  char- 
acterize tin.  Furthermore,  the  reactions  of  the  sulphides  of 
tin,  as  given  above,  are  of  importance,  as  is  also  the  fact  that 
the  hydroxides  are  precipitated  by  alkaline  carbonates,  and  do 


478  OUTLINES  OF   CHEMISTRY 

not  dissolve  in  excess  of  the  latter.  Zinc  precipitates  spongy 
tin  from  solutions  of  tin  salts.  This  may  then  be  dissolved  in 
hydrochloric  acid  and  tested  with  mercuric  chloride.  In  this 
way,  tin  may  be  distinguished  from  antimony. 

Occurrence,  Metallurgy,  and  Properties  of  Lead.  —  Lead, 
plumbum,  is  seldom  found  in  the  uncombined  state  in  nature. 
It  occurs  chiefly  as  the  sulphide  PbS,  which  commonly  crystal- 
lizes in  cubes  of  the  regular  system,  and  is  called  galenite. 
The  following  ores  are  also  found,  though  rarely  and  in  smaller 
quantities:  cerussite  PbCO3,  wulfenite  PbMoO4,  crocoisite 
PbCrO4,  bouronite  Cu2S  •  Sb2S3  -  2  PbS,  anglesite  PbSO4,  pyro- 
morphite  PbCl2  •  Pb3(PO4)2.  Galenite,  or  galena,  is  found  in 
fairly  large  quantities  in  the  United  States,  Great  Britain,  Ger- 
many, Spain,  and  Australia. 

To  obtain  lead  from  galena,  the  latter  is  roasted  so  as  to 
oxidize  it  in  part  to  oxide  and  in  part  to  sulphate,  thus  :  — 

2  PbS  -I-  3  O2  =  2  PbO  +  2  SO2,  and 

13 ua    i    o  r\         DT~Q/"V 


PbS  +  2  02  =  PbS04. 


By  then  strongly  heating  this  mixture  of  lead  oxide,  lead  sul- 
phate, and  unchanged  lead  sulphide,  out  of  contact  with  the 
air,  lead  is  obtained,  thus  :  — 

2  PbO  +  PbS  =  SO2  +  3  Pb,  and 
PbS04  +  PbS  =  2  S02  +  2  Pb. 

Lead  may  also  be  prepared  from  galena  by  heating  the  latter 
with  iron,  thus  :  — 

PbS  +  Fe  =  FeS  +  Pb. 

As  the  ferrous  sulphide  formed  is  much  lighter  than  lead,  it 
floats  on  top  of  the  latter,  and  is  run  off  like  a  slag.  The  mol- 
ten lead  can  readily  be  drawn  off  from  below.  The  separation 
of  lead  from  silver,  which  commonly  accompanies  lead,  has 
already  been  described.  The  world's  annual  production  of 
lead  is  about  950,000  tons.  In  the  year  1910  the  United 
States  furnished  372,227  tons. 

Lead  is  a  soft,  malleable,  bluish  white  metal,  having  a  bright 
luster  on  freshly  cut  surfaces,  which  soon  becomes  dull  on  ex- 
posure to  the  air  because  of  the  formation  of  a  film  of  oxide. 
The  specific  gravity  of  lead  is  11.4.  It  melts  at  327°,  and 


LEAD   AND   TIN  479 

boils  at  about  1200°  in  vacuo.  By  heating  lead  in  the  air,  the 
oxide  is  formed.  Water  in  contact  with  the  air  acts  on  lead 
somewhat,  forming  lead  hydroxide,  which  is  slightly  soluble. 
Waters  containing  calcium  sulphate  or  bicarbonate  act  on  lead, 
forming  lead  sulphate  or  carbonate.  These  salts  are  insoluble, 
and  so  form  a  coating  that  protects  the  metal  from  further 
action.  Hence  it  is  quite  feasible  to  use  lead  pipes  for  conduct- 
ing drinking  water.  Hydrochloric  and  sulphuric  acids  act  on 
lead  but  slightly,  because  the  resulting  products,  lead  chloride 
and  lead  sulphate,  are  sparingly  soluble.  In  nitric  or  acetic 
acid  lead  is  readily  soluble.  Many  other  weak  organic  acids 
dissolve  lead,  hence  it  is  unsuitable  for  cooking  utensils.  The 
atomic  weight  of  lead  is  207.10,  and  its  valence  is  generally 
two,  though  in  some  compounds  the  element  is  quadrivalent. 
From  solutions  of  its  salts,  lead  is  displaced  by  zinc,  iron,  and 
tin.  When  a  strip  of  zinc  is  hung  in  a  dilute  lead  acetate  so- 
lution, a  bulky,  branching  mass  of  lead,  known  as  a  lead  tree, 
is  formed  thus  :  — 

Zn  +  Pb(C2H302)2  =  Zn(C2H302)2  +  Pb. 

Uses  of  Lead.  —  Lead  has  been  known  since  earliest  times. 
The  Romans  used  it  for  water  pipes,  and  Pliny  distinguished 
between  lead  and  tin.  Much  lead  is  used  as  lead  pipes  in 
plumbing.  These  pipes  are  readily  made  of  hot,  plastic  lead 
by  means  of  hydraulic  pressure.  In  the  sulphuric  acid  indus- 
try, lead  is  used  for  lining  the  chambers  and  making  evaporat- 
ing dishes  and  pipes.  It  is  also  employed  in  other  chemical 
operations  for  containers.  Sheet  lead  further  serves  for  roofs 
and  gutters.  Alloyed  with  tin,  lead  forms  solder,  pewter,  and 
Britannia  metal,  which  have  already  been  described.  It  also 
enters  into  the  composition  of  Rose's  metal  and  Wood's 
metal,  which  are  alloys  of  low  melting  points.  Type  metal 
is  an  alloy  of  lead  with  antimony,  which  has  been  mentioned 
in  connection  with  the  latter.  Lead  used  for  shot  and  bullets 
contains  from  0.2  to  0.4  per  cent  of  arsenic.  Babbitt  metal 
consists  of  70  to  90  per  cent  lead  alloyed  with  tin  and  antimony. 
It  is  used  for  bearings  in  machines.  Large  quantities  of  lead 
are  also  used  in  making  storage  batteries.  Besides  this,  much 
lead  is  consumed  in  the  production  of  various  lead  compounds. 

Oxides   of  Lead.  —  The   following   five   oxides   of  lead   are 


480  OUTLINES   OF   CHEMISTRY 

known:  Pb2O,  PbO,  Pb2O3,  Pb3O4,  and  PbO2.  Lead  suboxide 
Pb2O  is  a  black  powder,  formed  by  the  action  of  the  air  on 
lead,  or  by  heating  lead  oxalate,  thus :  — 

2  PbC204  =  Pb20  +  3  C02  +  GO. 

Lead  oxide,  or  plumbic  oxide,  PbO  is  formed  by  burning  lead 
in  the  air.  It  results  as  a  by-product  in  the  separation  of  lead 
from  silver  by  cupellation.  Lead  oxide  is  also  made  by  calcin- 
ing the  carbonate  or  nitrate.  It  is  a  yellow,  amorphous  pow- 
der; but  when  fused  and  allowed  to  cool,  it  forms  a  crystalline 
mass.  This,  when  pulverized,  is  termed  litharge.  It  is  used  in 
making  glass  of  high  refracting  power,  in  glazing  pottery,  deco- 
rating porcelain,  and  preparing  many  lead  compounds. 

Lead  hydroxide  Pb(OH)2  is  formed  by  adding  caustic  alkali 
to  a  solution  of  a  lead  salt.  Lead  hydroxide  is  a  strong  base 
and  is  appreciably  soluble  in  water,  imparting  a  faintly  alkaline 
reaction  to  the  latter.  With  caustic  alkalies,  however,  lead 
hydroxide  exhibits  acid  properties,  forming  plumbites,  thus  :  — 

2  NaOH  +  Pb(OH)2  =  Na2PbO2  +  2  H2O. 

On  boiling  plumbic  oxide  with  caustic  alkalies,  plumbites  are 
also  formed. 

Lead  sesquioxide  Pb2O3  is  an  orange-yellow  powder,  obtained 
by  the  action  of  a  hypochlorite  on  a  solution  of  a  plumbite. 

Red  lead,  or  minium,  Pb3O4  is  formed  by  heating  plumbic 
oxide  in  the  air  to  about  400°.  It  is  a  bright  red  powder, 
which  is  generally  contaminated  with  some  litharge,  just  as  the 
latter  always  contains  small  amounts  of  minium.  Red  lead  is 
used  as  a  pigment  in  paints.  When  treated  with  nitric  acid,  it 
reacts  thus :  — 

Pb3O4  +  4  HNO8  =  2  Pb(NO3)2  +  PbO2  +  2  Had. 

It  is  evident,  then,  that  minium  behaves  like  a  mixture  of 
2  PbO  and  PbO2.  However,  it  is  probably  a  plumbate  of  lead, 
Pb2PbO4;  and  as  nitric  acid  acts  upon  it,  lead  nitrate  and 
plumbic  acid  H4PbO4  are  formed.  The  latter,  being  very 
unstable,  is  decomposed  at  once  into  2  H2O  and  PbO2. 

Lead  peroxide  PbO2  is  a  brown  powder,  which  is  also  formed 
by  the  action  of  hypochlorites  upon  lead  salts  in  alkaline  solu- 
tion, and  at  the  anode  during  the  electrolysis  of  lead  salts. 


LEAD   AND  TIN  481 

With  sulphuric  acid  it  forms  lead  sulphate  and  oxygen,  and 
with  hydrochloric  acid  lead  chloride  and  chlorine,  thus :  — 

2  Pb02  +  2  H2S04  =  2  PbS04  +  2  H2O  +  O2. 
PbO2  +  4  HC1  =  Pb012  +  2  H2O  +  C12. 

In  hot,  concentrated  solutions  of  caustic  alkalies  lead  peroxide 
dissolves,  forming  plumbates,  thus :  — 

2  KOH  +  Pb02  =  K2Pb03  +  H20. 

Plumbates  are  salts  of  metaplumbic  acid  H2PbO3.      They  are 
analogous  to  carbonates,  silicates,  and  stannates. 

All  oxides  of  lead  when  strongly  ignited  in  the  air  are  finally 
converted  into  litharge,  thus  :  — 

Pb2O  +O  =  2  PbO. 
Pb203          =  2  PbO  +  O. 
Pb304          =  3  PbO  -f  O. 
PbO2  =     PbO  +  O. 

Halides  of  Lead.  — Lead  chloride  PbCl2  is  a  white  salt  obtained 
as  a  precipitate  by  adding  a  soluble  chloride  to  a  solution  of  a 
lead  salt.  It  is  sparingly  soluble  in  cold  water,  but  100  parts 
of  boiling  water  dissolve  about  4  parts  of  the  salt.  On  ignition 
in  the  air,  lead  chloride  forms  lead  oxychloride  Pb2OCl2.  This 
oxychloride  is  also  produced  as  Pb2OCl2  •  H2O,  Pattinson's  white 
lead,  by  adding  milk  of  lime  to  a  hot  solution  of  lead  chloride :  — 

2  PbCl2  +  Ca(OH)2  =  Pb2OCl2  .  H2O  +  CaCl2. 

Lead  tetrachloride  PbCl4  is  a  yellow  oil  of  specific  gravity 
3.2,  which  congeals  as  a  crystalline  mass  at  — • 15°.  It  is  formed 
by  dissolving  lead  peroxide  in  well-chilled  concentrated  hydro- 
chloric acid,  or  by  passing  chlorine  into  a  cold  solution  of  lead 
chloride  in  hydrochloric  acid.  On  then  adding  ammonium 
chloride,  the  double  salt  PbCl4  •  2  NH4C1  is  obtained  as  a  crys- 
talline precipitate,  which  when  treated  with  concentrated  sul- 
phuric acid  at  0°  yields  lead  perchloride  as  a  yellow  oil.  The 
latter  readily  decomposes  into  PbCl2  and  C12.  At  105°  this 
decomposition  proceeds  with  explosive  violence. 

Lead  bromide  PbBr2  and  oxybromide  Pb2OBr2  are  analogous 
to  the  corresponding  chlorides. 

Lead  iodide  PbI2  is  a  yellow  precipitate  formed  by  adding  a 

2i 


482  OUTLINES  OF   CHEMISTRY 

soluble  iodide  to  a  solution  of  a  lead  salt.  It  dissolves  in  about 
200  parts  of  boiling  water,  from  which  it  crystallizes  in  shining 
hexagonal  scales.  Like  the  iodide  of  mercury,  it  readily  forms 
double  salts  with  iodides  of  the  alkalies. 

Lead  Nitrate  Pb(NO3)2  is  formed  by  the  action  of  nitric  acid 
on  lead,  lead  oxide,  or  lead  carbonate.  It  forms  octahedra 
which  are  soluble  in  about  2  parts  of  water. 

Lead  Acetate  Pb(C2H3O2)2  •  3  H2O  is  made  by  dissolving  lead 
or  lead  oxide  in  acetic  acid.  The  salt  forms  prismatic  crystals 
that  readily  effloresce.  They  are  very  soluble  in  water,  and 
the  solution  has  a  sweetish  taste,  hence  the  salt  is  called  sugar 
of  lead,  or  saccharum  saturni.  Lead  acetate  solutions  dissolve 
lead  oxide  or  hydroxide,  thus  forming  basic  lead  acetate 
Pb(C2H3O2)2  •  (PbO)^,  dilute  solutions  of  which  (about  2  per 
cent)  are  known  as  lead  water.  These  basic  lead  acetate  solu- 
tions readily  become  milky,  due  to  the  formation  of  lead  car- 
bonate with  carbon  dioxide  of  the  air. 

Lead  Sulphate  PbSO4  is  a  white  crystalline  precipitate  formed 
by  adding  a  soluble  sulphate  to  a  solution  of  a  lead  salt.  It  is 
practically  insoluble  in  water  and  dilute  sulphuric  acid.  But  it 
dissolves  appreciably  in  hydrochloric  and  nitric  acids,  and  fairly 
readily  in  concentrated  sulphuric  acid.  It  is  also  soluble  in 
caustic  alkalies,  in  sodium  thiosulphate,  and  in  ammonium 
acetate  and  other  ammonium  salts  of  organic  acids. 

Lead  Sulphide  PbS  is  found  in  nature  as  already  stated.  It 
is  formed  as  a  black  precipitate  when  hydrogen  sulphide  is 
passed  into  a  solution  of  a  lead  salt.  In  dilute  acids  it  is  in- 
soluble also  in  sulphides  of  the  alkalies;  but  when  boiled  with 
concentrated  hydrochloric  acid,  it  forms  lead  chloride  and 
hydrogen  sulphide.  When  boiled  with  dilute  nitric  acid,  lead 
sulphide  forms  lead  nitrate,  while  concentrated  nitric  acid 
oxidizes  it  to  sulphate.  The  latter  change  also  takes  place 
slowly  when  moist  lead  sulphide  is  exposed  to  the  air. 

Lead  Arsenate  Pb3(AsO4)2  is  formed  by  treating  a  lead  acetate 
solution  with  sodium  arsenate.  Lead  arsenate  is  a  white  pow- 
der that  is  but  sparingly  soluble  in  water.  Like  Paris  green, 
lead  arsenate  is  used  for  exterminating  potato  bugs  and  other 
insects  that  infest  cultivated  plants. 

Lead  Carbonate  PbCO3  is  found  in  nature,  as  already  men- 
tioned. It  is  formed  as  a  white  precipitate  by  adding  lead 


LEAD   AND   TIN  483 

nitrate  to  a  solution  of  ammonium  carbonate.  When  a  solution 
of  a  lead  salt  is  precipitated  with  normal  sodium  or  potassium 
carbonate,  basic  lead  carbonates  are  obtained  whose  composi- 
tion varies  with  the  concentration  and  temperature  of  the  solu- 
tions used. 

White  lead,  which  is  used  in  paints,  is  a  basic  lead  carbon- 
ate, commonly  of  the  composition  Pb(OH)2  •  2  PbCO3.  It  is 
formed  by  the  action  of  carbon  dioxide  on  basic  lead  acetate. 
Thenard's  method,  also  called  the  French  method,  consists  of 
passing  carbon  dioxide  into  a  solution  of  basic  lead  acetate 
formed  by  dissolving  lead  oxide  in  lead  acetate  solution.  Thus 
basic  lead  carbonate  is  precipitated,  while  neutral  lead  acetate 
remains  in  solution  and  can  be  used  over  again.  This  is  a  rapid, 
direct  process,  but  it  yields  a  crystalline  or  granular  product 
which  does  not  possess  the  desirable  covering  properties  of 
white  lead  produced  by  the  Dutch  process.  The  latter  has  been 
in  use  for  nearly  three  hundred  years,  and  consists  of  allowing 
the  vapors  of  vinegar  and  carbon  dioxide  to  act  slowly  upon 
large  surfaces  of  lead.  To  effect  this,  the  sheet  lead  is  loosely 
rolled  together  in  spirals,  each  of  which  is  placed  in  a  small 
earthenware  pot  into  which  vinegar,  4  to  5  per  cent  acetic  acid, 
has  been  poured.  The  lead  rests  on  a  shelf  so  as  not  to  be  in 
contact  with  the  vinegar.  Lead  gratings  or  other  forms  of 
lead  castings  are  sometimes  used  instead  of  sheet  lead  spirals. 
The  earthen  pots,  which  are  about  20  cm.  high  and  12  cm.  in 
diameter,  are  loosely  covered  and  set  in  horse  manure  or  spent 
tan  bark.  The  acetic  acid  thus  vaporizes  and  acts  upon  the 
lead,  forming  basic  lead  acetate,  which  is  then  acted  upon  by 
the  carbon  dioxide  liberated  during  the  process  of  fermentation 
of  the  manure  or  spent  tan,  resulting  in  the  formation  of  white 
lead.  Several  weeks  are  required  to  complete  the  action.  The 
fermentation  process  not  only  furnishes  the  required  carbon 
dioxide,  but  also  produces  the  heat  necessary  to  vaporize  the 
vinegar.  An  amorphous  white  lead  of  superior  covering  power 
is  obtained  in  this  way»  Many  other  processes  are  in  use,  and 
some  of  them  yield  white  lead  of  excellent  quality.  They  all 
depend  upon  the  action  of  acetic  acid  and  carbon  dioxide  upon 
lead.  Electrolytic  processes  have  also  been  proposed;  they  de- 
pend upon  the  fact  that  when  the  electric  current  is  passed  be- 
tween two  lead  plates  dipping  in  a  solution  of  sodium  nitrate 


484  OUTLINES   OF   CHEMISTRY 

or  chlorate,  and  sodium  bicarbonate,  basic  lead  carbonate  is 
precipitated.  It  is  claimed  that  the  product  obtained  is 
amorphous. 

White  lead  is  ground  with  linseed  oil  and  used  as  a  paint.  The 
oil  gives  the  paint  a  yellowish  tinge  which  is  usually  dispelled 
by  adding  a  trace  of  blue  or  black  pigment.  Barium  sulphate, 
lead  sulphate,  and  calcium  carbonate,  particularly  the  former, 
are  used  as  adulterants  of  white  lead.  Their  presence  can 
readily  be  detected.  White  lead  turns  dark  when  exposed  to 
hydrogen  sulphide,  or  when  sulphide  pigments  are  mixed  with 
it.  It  is  also  poisonous,  and  care  must  be  exercised  in  its  manu- 
facture and  use  so  that  it  will  not  get  into  the  system.  All 
lead  compounds  are  poisonous ;  the  readily  soluble  ones  are,  of 
course,  the  most  dangerous.  Painters  and  others  that  work 
with  lead  or  its  compounds  are  liable  to  lead  colic,  which  is 
produced  by  the  gradual  absorption  of  lead  compounds,  by 
breathing  the  latter  in  form  of  fine  dust,  or  by  constantly 
handling  them. 

Analytical  Tests  for  Lead.  —  When  mixed  with  soda  and 
heated  on  charcoal  before  the  blowpipe,  lead  compounds  yield 
globules  of  metallic  lead,  which  upon  oxidation  are  transformed 
to  plumbic  oxide.  Lead  sulphide,  sulphate,  hydroxide,  chlo- 
ride, and  carbonate  are  all  characteristic  precipitates  which  may 
be  obtained  from  solutions  of  lead  salts  as  already  stated.  A 
lemon-yellow  precipitate  of  lead  chromate  is  produced  by  add- 
ing potassium  chromate  or  bichromate  to  a  solution  of  a  lead 
salt. 


CHAPTER   XXVIII 

CHROMIUM,    MOLYBDENUM,    TUNGSTEN,    AND    URANIUM 

CHROMIUM  (Cr  —  52.0),  molybdenum  (Mo  —  96.0),  tung- 
sten (W  — 184.0),  and  uranium  (U  —  238.5)  are  related  to  the 
members  of  the  sulphur  family  about  as  titanium,  zirconium, 
cerium,  and  thorium  are  to  the  carbon  and  silicon  group,  or  as 
vanadium,  columbium,  and  tantalum  are  to  the  phosphorus 
group.  The  members  of  the  chromium  family  form  the  tri- 
oxides  CrO3,  MoO3,  WO3,  and  UO3,  which  are  analogous  to 
SO3.  Like  the  latter,  they  are  acid  anhydrides.  With  me- 
tallic oxides  they  form  salts  like  K2CrO4,  K2MoO4,  K2WO4, 
which  correspond  to  K2SO4.  Again,  the  compounds  MoO2, 
WO2,  and  UO2  are  analogous  to  SO2.  On  the  other  hand, 
chromium  often  acts  as  a  base,  forming  compounds  that  are 
analogous  to  those  of  aluminum  and  iron,  as  in  CrClg,  and 
Cr2(SO4)3.  The  other  members  of  the  group  do  not  form 
such  salts,  but  they  enter  into  other  rather  complicated  com- 
pounds. This  is  particularly  the  case  with  uranium.  Chro- 
mium is  the  most  important  member  of  the  group.  None  of  the 
elements  of  this  family  occur  in  nature  in  the  free  state. 

Occurrence,  Preparation,  and  Properties  of  Chromium.  —  Chro- 
mium is  generally  found  in  form  of  chromite,  also  called  chrome 
iron  ore,  FeO  •  Cr2O3  or  Fe(CrO2)2,  which  crystallizes  in  octa- 
hedra  and  is  isomorphous  with  spinel.  Crocoisite  PbCrO4,  first 
found  in  Siberia,  occurs  more  rarely,  though  it  was  in  this 
mineral  that  chromium  was  discovered  in  1797  by  Vauquelin. 
The  name  chromium  was  given  the  element  because  it  forms 
colored  compounds. 

Metallic  chromium  may  readily  be  prepared  by  Goldschmidt's 
process,  consisting  of  igniting  a  mixture  of  chromic  oxide  and 
finely  divided  aluminum  by  means  of  a  fuse  of  magnesium  rib- 
bon or  a  mixture  of  barium  peroxide  and  aluminum  powder. 
The  reaction  when  thus  started  proceeds  to  completion.  Chro- 
mium may  also»be  obtained  by  reducing  the  oxide  with  carbon 
in  the  electric  furnace. 

485 


486  OUTLINES  OF  CHEMISTRY 

Chromium  is  a  steel-gray,  hard,  brittle  metal  of  high  metallic 
luster.  Its  specific  gravity  is  6.8.  It  requires  the  electric 
furnace  to  melt  chromium.  The  metal  may  be  polished,  and 
it  remains  unchanged  in  the  air.  At  high  temperatures  it 
burns  in  oxygen  or  in  the  air,  emitting  a  brilliant  light  and 
forming  Cr2O3.  Chromium  is  not  attacked  by  nitric  acid; 
but  in  warm  dilute  hydrochloric  or  sulphuric  acid  it  dissolves 
with  evolution  of  hydrogen.  Chromium  is  commonly  bivalent, 
trivalent,  or  hexavalent. 

Chromium  is  used  in  the  steel  industry  for  making  chrome  steel, 
in  which  process  ferrochromium,  an  alloy  of  iron  and  chromium 
containing  60  per  cent  of  the  latter,  is  added  to  the  steel.  This 
ferrochromium  is  readily  prepared  by  heating  chromium  ore 
with  carbon  in  the  electric  furnace. 

Chromic  Oxide  and  Hydroxides.  —  With  oxygen,  chromium 
forms  a  sesquioxide,  Cr2O3,  which  acts  as  a  base,  and  a  trioxide, 
CrO3,  which  acts  as  an  acid.  These  are  the  only  oxides  of 
chromium  that  are  known  with  certainty,  though  chromous 
hydroxide  Cr(OH)2,  in  which  the  metal  is  bivalent,  has  also 
been  prepared,  being  formed  as  a  yellow  precipitate  when 
caustic  alkali  is  added  to  a  solution  of  chromous  chloride. 

Chromic  oxide  (chromium  sesquioxide)  Cr2O3  is  a  grass-green 
powder  formed  by  ignition  of  chromic  hydroxide,  chromium 
tripxide,  or  ammonium  bichromate.  The  oxide  may  be  obtained 
in  the  form  of  very  dark  green,  lustrous,  hexagonal  crystals  by 
passing  the  vapors  of  chromyl  chloride  CrO2Cl2  through  a  red- 
hot  tube,  thus :  — 

2Cr02Cl2=Cr203  +  2Cl2  +  0. 

Amorphous  chromic  oxide  dissolves  readily  in  acids,  but  after 
strong  ignition  the  latter  scarcely  attack  it ;  and  like  highly 
heated  oxides  of  iron  or  aluminum,  chromic  oxide  is  then  usually 
fused  with  bisulphate  of  potassium  in  order  to  effect  its  solu- 
tion. Chromic  oxide  is  used  as  a  pigment,  chrome  green,  in 
paints.  It  also  serves  in  coloring  glass  green. 

Chromic  hydroxide  is  formed  as  a  grayish  blue  precipitate  of 
the  composition  Cr(OH)3  •  2  H2O  when  ammonia  is  added  to  a 
solution  of  a  chromium  salt.  On  drying  this  substance  over 
sulphuric  acid  in  a  vacuum,  a  residue  of  very  nearly  the  com- 
position Cr(OH)3  may  be  obtained.  On  heating  the  latter  in 


CHROMIUM,    MOLYBDENUM,   TUNGSTEN,   AND  URANIUM      487 

hydrogen  to  about  220°,  chromous  hydroxide  Cr(OH)2  is 
formed,  which  upon  strong  ignition  yields  chromic  oxide  and 
water.  Potassium  or  sodium  hydroxide  also  precipitates 
hydrated  chromic  hydroxide  from  solutions  of  chromium  salts, 
but  the  precipitate  always  contains  some  alkali.  In  excess  of 
potassium  or  sodium  hydroxide,  chromic  hydroxide  is  soluble, 
forming  chromites,  thus :  — 

NaOH  +  Cr(OH)3  =  NaCrO2  +  2  H2O. 

This  behavior  is  similar  to  that  of  aluminum  hydroxide.  How- 
ever, on  boiling  solutions  of  chromites,  they  are  decomposed, 
chromic  hydroxide  being  precipitated,  while  aluminates  are 
stable  under  similar  treatment.  Insoluble  chromites  are  also 
known,  chrome  iron  ore  Fe(CrO2)2  being  a  compound  of  this 
character.  With  acids,  chromic  hydroxide  forms  chromic 
salts ;  indeed,  it  generally  acts  as  a  base. 

Chromous  Compounds. — In  these  compounds  chromium  is 
bivalent.  The  hydroxide,  Cr(OH)2,  has  already  been  men- 
tioned. Chromous  chloride  CrCl2  and  chromous  sulphate 
CrSO4  •  7  H2O  have  been  prepared  by  dissolving  the  metal  in 
hydrochloric  or  sulphuric  acid.  Chromous  chloride  has  also 
been  made  by  heating  chromic  chloride  in  hydrogen.  Chro- 
mous salts  have  been  studied  but  little.  They  are  strong 
reducing  agents,  for  they  readily  pass  over  into  chromic  com- 
pounds. 

Chromic  Salts.  —  These  are  made  by  action  of  acids  on  chro- 
mic hydroxide.  Chromic  chloride  CrCl3  may  also  be  prepared 
by  the  action  of  chlorine  on  heated  chromium,  or  by  passing 
chlorine  over  a  red-hot  mixture  of  chromic  oxide  and  carbon. 
Thus  prepared,  the  compound  sublimes  in  violet  leaflets  that 
are  almost  insoluble  in  water.  By  ignition  in  the  air  these 
pass  over  into  chromic  oxide.  On  long-continued  boiling  the 
salt  slowly  dissolves ;  but  when  chromous  chloride  is  added, 
even  in  traces,  the  action  progresses  much  more  rapidly,  a 
green  solution  being  formed  from  which  by  evaporation  green, 
deliquescent  crystals,  CrCl3  •  6  H2O,  may  be  obtained.  The 
latter  yield  both  water  and  hydrochloric  acid  on  being  heated 
in  the  air;  consequently,  to  prepare  the  anhydrous  chloride 
from  them,  they  are  heated  in  a  current  of  chlorine  or  hydro- 
chloric acid  gas. 


488  OUTLINES   OF   CHEMISTRY 

Chromic  sulphate  Cr2(SO4)3  •  15  H2O  is  deposited  from  cold 
solutions  in  the  form  of  violet  crystals.  The  salt  forms  violet 
solutions  at  ordinary  temperatures.  On  boiling  these,  they 
become  green  because  a  hydrolysis  takes  place,  the  exact  nature 
of  which  has  not  yet  been  definitely  determined.  From  the 
green  solutions  no  crystals  are  obtainable ;  but  on  standing, 
these  solutions  slowly  become  violet  again  and  deposit  violet 
crystals.  This  peculiar  behavior  is  also  exhibited  by  other 
chromic  salts. 

Chrome  alums  are  double  salts  which  chromic  sulphate  forms 
with  sulphates  of  the  alkalies.  They  have  the  same  general 
formula  as  other  alums,  and  are  isomorphous  with  them.  Potas- 
sium chrome  alum  KCr(SO4)2  •  12  H2O  is  the  commonest  of 
these  compounds,  and  is  generally  called  simply  chrome  alum. 
It  forms  octahedra  which  are  of  a  dark  violet  color,  but  appear 
reddish  by  transmitted  light.  They  effloresce  on  exposure  to 
the  air.  Their  solutions  exhibit  color  changes  similar  to  those 
of  chromic  sulphate. 

Chromates,  Bichromates,  and  Chromium  Trioxide.  —  The  chief 
source  of  chromium  is  chrome  iron  ore,  as  already  stated.  When 
this  is  pulverized,  mixed  with  potash  and  calcium  carbonate, 
and  roasted  in  contact  with  air,  ferric  oxide,  potassium  chro- 
mate,  calcium  chromate,  and  carbon  dioxide  are  formed,  thus  :  — 

4  Fe(Cr02)2  +  6  K2CO3  +  2  CaCO3  +  7  O2 

=  6  K2CrO4  +  2  CaCrO4  +  2  Fe2O3  +  8  CO2. 

On  lixiviating  the  mass  with  water,  potassium  and  calcium 
chromates  dissolve,  and  on  adding  potassium  sulphate  to  the 
solution,  calcium  sulphate  is  formed,  thus :  — 

CaCr04  +  K2S04  =  CaSO4  +  K2CrO4. 

On  evaporating  the  clear  solution  of  potassium  chromate  thus 
obtained,  the  salt  crystallizes  out  in  rhombic  pyramids  of  lemon 
color  which  are  isomorphous  with  potassium  sulphate.  Potas- 
sium chromate  is  soluble  in  about  2  parts  of  water.  The  solu- 
tion has  an  alkaline  reaction  because  of  partial  hydrolysis  of 
the  salt. 

On  adding  sulphuric  acid  to  a  solution  of  potassium  chro- 
mate, the  latter  turns  orange-red  and  readily  deposits  large,  red, 
triclinic  crystals  of  potassium  bichromate  K2Cr2O7.  These  are 


CHROMIUM,  MOLYBDENUM,  TUNGSTEN,  AND  URANIUM   489 

soluble  in  about  1  part  of  water  at  100°,  whereas  at  room  tem- 
peratures about  10  parts  of  water  are  necessary  to  effect  their 
solution.  The  salt  can  consequently  readily  be  purified  by 
recrystallization.  Potassium  bichromate  is  analogous  to  po- 
tassium pyrosulphate  K2S2O7.  On  heating  potassium  bichro- 
mate, it  melts  and  then  decomposes  :  — 

2  K2Cr2O7  =  2  K2CrO4  +  CraO8  +  3  O. 

When  heated  with  concentrated  sulphuric  acid,  chrome  alum 
and  oxygen  are  formed  :  — 

K2Cr207  +  4  H2SO4  =  2  KCr(SO4)2  +  4  H2O  +  3  O. 

With  hot  concentrated  hydrochloric  acid,  chlorine  is  evolved :  — 

K2Cr207  +  14  HC1  =  2  CrCl3  +  2  KC1  +  7  H2O  +  3  C12. 

From  these  reactions  it  is  clear  that  potassium  bichromate  is  a 
strong  oxidizing  agent,  each  molecule  yielding  three  oxygen  atoms 
that  are  available  to  effect  oxidations. 

If,  in  the  roasting  of  chrome  iron  ore  with  calcium  carbonate 
and  potash,  soda  is  substituted  for  the  latter,  sodium  chromate 
is  produced.  It  is  a  yellow  salt  that  may  be  obtained  in  deli- 
quescent, prismatic  crystals,  Na2CrO4  •  10  H2O,  which  are  iso- 
morphous  with  Glauber's  salt.  On  treatment  with  sulphuric 
acid,  sodium  chromate  may  be  transformed  to  sodium  bichro- 
mate Na2Cr2O7  •  2  H2O,  which  forms  red,  triclinic  crystals, 
soluble  in  about  1  part  of  water  at  ordinary  temperatures. 
This  salt  is  cheaper  than  potassium  bichromate  and  hence  is 
generally  used  in  place  of  the  latter  in  the  industries.  By 
adding  potassium  chloride  to  solutions  of  sodium  bichromate, 
potassium  bichromate  may  be  obtained,  for  it  is  less  soluble. 

Gelatine  treated  with  potassium  bichromate  solution  darkens 
on  exposure  to  light  and  becomes  insoluble,  probably  because 
of  the  formation  of  chromic  oxide  and  partial  oxidation  of  the 
gelatine.  This  fact  is  used  in  photographic  processes,  and  also 
in  making  a  glue  that  gradually  becomes  insoluble.  Bichro- 
mates are  used  in  dyeing  and  tanning,  also  as  oxidizing  agents  in 
the  laboratory  and  in  the  industries.  Ammonium  bichromate 
(NH4)2Cr2O7  forms  readily  soluble  red  crystals  which  on  igni- 
tion continue  to  oxidize  with  brilliant  scintillations  :  — 

(NH4)2Cr2O7  =  Cr2O3  +  4  H2O  +  Na. 


490  OUTLINES   OF  CHEMISTRY 

Lead  chromate  PbCrO4  is  a  bright  yellow  precipitate  formed 
by  adding  either  a  soluble  chromate  or  bichromate  to  a  solution 
of  a  lead  salt.  It  serves  as  a  pigment  in  paints,  being  called 
chrome -yellow.  It  may  be  fused  without  decomposition.  On 
cooling,  it  then  forms  a  brown  crystalline  solid,  which  when 
heated  with  carbon  compounds  readily  oxidizes  the  latter,  hence 
its  use  in  organic  combustion  analyses. 

Barium  chromate  BaCrO4  is  a  yellow  precipitate  prepared 
by  adding  a  soluble  chromate  or  bichromate  to  a  solution  of  a 
barium  salt.  It  is  insoluble  in  acetic  acid,  but  soluble  in  nitric 
or  hydrochloric  acid.  It  is  also  used  as  a  pigment,  under  the 
name  ultramarine  yellow.  Calcium  chromate  CaCrO4  •  2  H2O 
is  analogous  to  gypsum  CaSO4  •  2  H2O.  Magnesium  chromate 
MgCrO4  •  7  H2O  is  analogous  to  Epsom  salt  MgSO4  •  7  H2O. 
With  the  exception  of  barium  chromate,  the  chromates  of  the  alka- 
line earths  all  dissolve  in  acetic  acid,  which  fact  is  used  in  analysis. 
Chromates  of  the  heavy  metals  are  insoluble  in  water  and  hence 
are  readily  prepared  by  precipitation.  Silver  chromate  Ag2CrO4 
and  mercurous  chromate  Hg2CrO4  are  characteristic  red  pre- 
cipitates. 

On  adding  concentrated  sulphuric  acid  to  a  well-cooled  con- 
centrated solution  of  sodium  or  potassium  bichromate,  chro- 
mium trioxide,  or  chromic  acid  anhydride,  CrO8  separates  out  in 
form  of  beautiful,  dark  red,  deliquescent,  rhombic  needles.  Its 
aqueous  solutions  probably  contain  bichromic  acid  H2Cr2O7,  but 
the  latter  has  not  been  isolated.  On  neutralization  with  alka- 
lies, such  solutions  of  chromic  acid  yield  bichromates  and  chro- 
mates. Chromium  trioxide  is  a  powerful  oxidizing  agent.  It 
destroys  organic  tissues  and  oxidizes  many  compounds,  thus :  — 

2  CrO3  +  3  H2S  =  Cr2O3  +  3  S  +  3  H2O. 
CrO3  +  6  HC1  =  CrCl3  +  3  H2O  +  3  01. 
2Cr03  +  3S02=Cr2(S04)3. 
2  CrO8  +  3  H2SO4  +  3  C2H5 .  OH 

=  Cr2(S04)3  +  3  CH3CHO  +  6  H2O. 
2  Cr08  -f  3  H2S04  +  3  (COOH)2  =  Cr2(SO4)3  +  6  H2O  +  6  CO2. 

It  should  be  borne  in  mind  that  compounds  of  chromium  are 
poisonous,  those  that  are  readily  soluble  being  most  harmful. 
Chromyl  Chloride  CrO2  •  C12  is  analogous  to  sulphuryl  chloride 


CHROMIUM,    MOLYBDENUM,    TUNGSTEN,   AND   URANIUM      491 

SO2  •  C12.  It  is  a  dark  red,  fuming  liquid  commonly  made  by 
distilling  a  mixture  of  sodium  chloride,  sodium  or  potassium 
bichromate,  and  sulphuric  acid,  thus:  — 

4  NaCl  +  Na2Cr2O7+6  H2SO4  =  6  NaHSO4  +  3  H2O  + 2  CrO2Cl2. 
Water  decomposes  the  compound:  — 

Cr02  •  C12  +  H20  =  Cr03  +  2  HC1 ; 

hence,  in  preparing  it,  a  sufficient  amount  of  sulphuric  acid 
must  be  used  to  absorb  the  water  formed  during  the  reaction. 

Analytical  Tests  for  Chromium  Compounds.  —  Because  of  their 
color,  these  compounds  are  readily  detected.  Chromous  com- 
pounds readily  pass  over  into  chromic  compounds.  The  latter 
yield  a  green  coloration  when  heated  in  the  borax  or  sodium 
metaphosphate  bead.  Caustic  alkalies  precipitate  chromic  hy- 
droxide, which  is  soluble  in  an  excess  of  the  precipitant,  but  is 
again  precipitated  on  boiling.  Alkaline  carbonates  precipitate 
chromic  hydroxide.  Hydrogen  sulphide  produces  no  precipi- 
tate in  solutions  of  chromic  salts,  but  ammonium  sulphide  pre- 
cipitates chromic  hydroxide.  The  latter,  like  the  hydroxide 
of  aluminum,  is  scarcely  affected  by  an  excess  of  ammonium 
hydroxide.  When  fused  with  soda  and  saltpeter,  chromic 
compounds  yield  chromates;  these  are  characterized  by  their 
yellow  color  and  by  the  characteristic  insoluble  precipitates 
which  their  solutions  yield  with  silver  nitrate,  barium  chloride, 
lead  nitrate,  and  mercurous  nitrate.  Furthermore,  when  acidi- 
fied with  sulphuric  acid,  chromate  solutions  readily  oxidize 
oxalic  acid,  alcohol,  ferrous  salts,  etc.,  yielding  green  solutions 
of  chromic  sulphate. 

Molybdenum. — This  metal  is  found  in  molybdenite  MoS2, 
which  looks  like  graphite,  and  in  wulfenite  PbMoO4,  which 
forms  yellowish  tetragonal  crystals.  On  roasting  molybden- 
ite, it  is  converted  to  molybdenum  trioxide  MoO3,  a  white  crys- 
talline powder  which  is  the  most  stable  oxide  of  molybdenum. 
This  powder  readily  dissolves  in  caustic  alkalies  or  ammonium 
hydroxide,  forming  molybdates.  Of  these  ammonium  molyb- 
date  (NH4)2MoO4  is  an  important  reagent  in  analytical  chemistry, 
for  with  phosphoric  acid  or  phosphates  in  dilute  nitric  acid  solu- 
tion it  yields  a  characteristic  light  yellow  precipitate  of  ammo- 
nium phosphomolybdate  (NH4)3PO4 . 11  MoO3  •  6  H2O.  The 


492  OUTLINES   OF   CHEMISTRY 

composition  of  the  latter  varies  somewhat  according  to  the 
conditions  of  precipitation.  In  dilute  acids  the  compound  is 
insoluble,  but  in  alkalies  or  an  excess  of  phosphoric  acid  it  dis- 
solves. On  treating  a  solution  of  a  molybdate  with  nitric  acid, 
yellow  crystals  of  molybdic  acid  H2MoO4  •  H2O  separate  out, 
which  upon  drying  yield  H2MoO4.  Similarly  when  ammo- 
nium phosphomolybdate  is  treated  with  aqua  regia,  crystals  of 
phosphomolybdic  acid  H3PO4 -HMoO3- 12  H2O  are  formed. 
This  acid  forms  insoluble  salts  with  salts  of  potassium,  ammo- 
nium, and  the  alkaloids,  and  is  consequently  of  value  in  ana- 
lytical tests  for  the  latter.  It  is  known  as  Sonnenschein's 
reagent. 

Molybdenum  receives  its  name  from  the  Greek  word  mean- 
ing lead,  for  molybdenite  was  regarded  as  graphite,  which  was 
called  black  lead.  Scheele  made  molybdic  acid  in  1778  by 
treating  molybdenite  with  nitric  acid,  thus  showing  that  the 
mineral  is  not  graphite.  In  1790  Hjelm  prepared  the  metal 
by  heating  the  oxides  or  chlorides  in  a  current  of  hydrogen, 
and  Klaproth  determined  the  true  nature  of  wulfenite  in  1797. 
Molybdenum  may  also  be  made  by  heating  the  oxide  with  car- 
bon in  the  electric  furnace.  It  is  a  hard,  lustrous  metal  of 
specific  gravity  9.1.  Its  physical  properties  resemble  those  of 
iron.  Like  the  latter,  it  may  be  welded  and  tempered.  Besides 
the  compounds  already  mentioned,  molybdenum  forms  the 
oxides  Mo2O3  and  MoO2;  the  chlorides  (MoCl2)3,  MoCl3, 
MoCl4,  MoCl5;  and  the  sulphides  MoS3  and  MoS4.  It  also 
forms  a  large  variety  of  additional,  rather  complicated,  com- 
pounds. 

Tungsten.  —  This  metal  is  found  in  the  minerals  scheelite 
CaWO4,  wolframite  (FeMn)WO4,  and  stolzite  PbWO4.  By 
treating  the  pulverized  ores  with  nitric  acid,  tungsten  trioxide 
WO3  is  obtained  as  a  yellow  powder  which  is  insoluble  in 
acids.  With  caustic  alkalies,  it  forms  tungstates  like  Na2WO4 
and  K2WO4,  which  are  soluble,  and  from  whose  solutions  tung- 
stic  acid  H2WO4  •  H2O  is  precipitated  by  means  of  acids. 
Sodium  tungstate  may  also  be  made  by  fusing  native  tungstates 
with  soda  and  lixiviating  the  mass  with  water.  The  salts 
Na2WO4  •  2  H2O  and  Na2W4O13  . 10  H2O  are  used  as  mordants. 
They  are  also  employed  in  making  fabrics  fireproof.  Tung- 
sten forms  the  oxides  WO2  and  WOC;  the  chlorides  WC12, 


CHROMIUM,   MOLYBDENUM,   TUNGSTEN,   AND   URANIUM       493 

WC14,  WC15,  WC16,  WO2C12,  and  WOC14;  and  the  carbides 
W2C  and  WC.  In  addition,  tungsten  enters  into  a  very  large 
number  of  complex  compounds.  Phosphotungstic  acid  is  known 
as  Scheibler's  reagent.  It  consists  of  a  compound  of  tungstic 
and  phosphoric  acids,  which  is  analogous  to  phosphomolybdic 
acid. 

Tungsten  was  discovered  in  1781  by  Scheele,  and  two  years 
later  the  element  was  isolated  by  the  d'Elhujar  brothers. 
Tungsten  may  be  made  by  igniting  the  oxides  or  chlorides  in 
a  current  of  hydrogen,  by  heating  the  oxides  with  carbon  in 
the  electric  furnace,  or  by  the  Goldschmidt  process.  It  is  a 
hard,  steel-gray,  lustrous,  brittle  metal  of  specific  gravity  19.13. 
Its  melting  point  lies  very  high.  On  ignition,  it  burns  to  the 
trioxide.  Tungsten  is  used  in  the  steel  industry,  for  its  pres- 
ence to  the  extent  of  5  to  8  per  cent  in  steel  makes  the  latter 
very  hard  and  tough.  Tungsten  is  now  also  being  used  in 
making  filaments  for  incandescent  electric  lamps,  which  have  a 
higher  efficiency  than  the  lamps  using  a  carbon  filament. 

Uranium.  —  The  mineral  pitchblende  or  uraninite  UO2-2UO3 
or  U3O8  constitutes  the  chief  source  of  uranium,  which,  like 
molybdenum  and  tungsten,  belongs  to  the  rarer  elements.  The 
true  nature  of  pitchblende  was  first  recognized  by  Klaproth  in 
1789;  but  metallic  uranium  was  not  prepared  till  1841,  when 
Peligot  obtained  it  by  heating  metallic  sodium  and  urarious 
chloride  together.  The  element  is  named  after  the  planet 
Uranus,  discovered  in  1787.  Metallic  uranium  may  also  be 
obtained  by  reduction  of  its  oxides  with  carbon  in  the  electric 
furnace,  by  heating  the  oxides  with  aluminum,  or  by  elec- 
trolysis of  the  molten  double  chloride  UC14  •  2  NaCl.  Uranium 
is  a  very  hard,  white,  lustrous  metal,  not  unlike  iron  in  appear- 
ance. Its  specific  gravity  is  18.7;  and  its  melting  point  is 
about  1500°.  It  has  the  highest  atomic  weight  of  all  the  ele- 
ments known,  namely,  238.5.  Uranium  shows  variable  valence 
in  its  compounds,  which  are  very  numerous.  The  maximum 
valence  exhibited  by  uranium  is  eight. 

On  treating  pitchblende  with  nitric  acid,  uranyl  nitrate 
UO2  •  (NO3)2  •  6  H2O  may  be  obtained.  This  salt  crystallizes  in 
greenish  yellow,  rhombic  prisms  that  show  fluorescence,  a 
characteristic  of  many  uranium  salts.  By  careful  ignition, 
uranyl  nitrate  may  be  converted  into  uranium  trioxide  UO8, 


494  OUTLINES   OF   CHEMISTRY 

a  dark  yellow  powder,  which  is  uranic  acid  anhydride.  On 
digesting  uranium  trioxide  with  nitric  acid,  uranic  acid 
UO2(OH)2  or  H2UO4  may  be  obtained  as  an  amorphous, 
yellow  powder.  Towards  strong  acids,  this  acts  as  a  base, 
forming  the  uranyl  salts  containing  the  bivalent  radical 
uranyl  UO2.  Thus  we  have  uranyl  chloride  UO2-C12;  uranyl 
sulphate  UO2  -  SO4  •  3  H2O  ;  uranyl  nitrate  UO2(NO3)2  •  6  H2O  ; 
uranyl  acetate  UO2(C2H3O2)2 ;  uranyl  ammonium  carbon- 
ate UO2.CO3-  2(NH4)2CO3;  uranyl  ammonium  phosphate 
UO2-NH4-PO4.  The  latter  salt  is  obtained  as  a  precipitate, 
insoluble  in  acetic  acid,  when  uranyl  acetate  is  added  to  a  solu- 
tion of  a  soluble  phosphate  containing  ammonium  chloride. 
This  fact  is  sometimes  used  in  the  estimation  of  phosphoric 
acid. 

By  the  addition  of  caustic  alkalies  to  solutions  of  uranyl 
compounds,  uranates  are  precipitated.  These  are  not  deriva- 
tives of  H2UO4,  but  of  H2U2O7;  that  is,  they  are  diuranates, 
analogous  to  pyrosulphates.  So  there  are  potassium  diuranate 
K2U2O7-3H2O  and  sodium  diuranate  Na2U2O7  •  6  H2O.  The 
latter  is  called  uranium  yellow,  and  is  used  in  making  a  beautiful 
yellowish  green,  fluorescent  uranium  glass. 

Besides  the  uranyl  salts,  uranium  forms  a  series  of  uranous 
compounds  in  which  the  element  is  quadrivalent.  So  there  are 
uranous  chloride  UC14,  uranous  sulphate  U(SO4)2  •  8  H2O,  uranous 
hydroxide  U(OH)4,  and  uranous  oxide  UO2.  The  latter  may 
be  obtained  as  a  steel-gray  crystalline  powder  by  ignition  of 
the  trioxide  in  hydrogen;  it  was  regarded  as  the  element  itself, 
till  Peligot  showed  that  this  was  erroneous.  Uranous  oxide  is 
used  in  making  a  fine  jet-black  glass. 

Uranium  Carbide  U2C3  may  be  made  by  heating  uranium 
oxides  with  carbon  in  the  electric  furnace.  It  is  harder  than 
quartz.  The  sulphides,  US2  and  UO2S,  are  also  known.  The 
variable  valence  of  uranium  is  readily  apparent  from  the  follow- 
ing list  of  its  chlorides  and  oxides  :  UC13,  UC14,  UC15;  UO2, 
U308,  U03,  U04. 

Compounds  of  uranium  exhibit  the  phenomena  of  radio- 
activity which  led  to  the  discovery  of  radium  as  already  stated. 
The  radio-activity  of  uranium  bearing  minerals  is  proportional  to 
their  uranium  content. 


CHAPTER  XXIX 

MANGANESE 

MANGANESE  is  a  metal  which  forms  quite  a  variety  of  com- 
pounds. The  element  is  bivalent  and  basic  in  character  in  the 
manganous  compounds  which  resemble  those  of  the  magnesium 
and  iron  groups.  In  manganic  compounds  manganese  is  trivalent, 
thus  resembling  aluminum,  chromium,  and  trivalent  iron.  In 
its  dioxide  and  the  manganites,  manganese  is  quadrivalent  and 
shows  analogies  to  tin  and  lead.  Again,  manganese  exhibits 
similarities  to  sulphur  and  the  chromium  group  by  forming  an 
acidic  trioxide  and  manganates  that  are  analogous  to  sulphates, 
chromates,  molybdates,  etc.  Finally,  the  metal  exhibits  re- 
semblances to  the  halogens  in  forming  a  heptoxide  and  per- 
manganic acid  and  permanganates,  which  are  analogous  to 
chlorine  heptoxide,  perchloric  acid,  and  the  perchlorates.  In 
the  periodic  system  of  the  elements  manganese  occurs  in  group 
VII  with  the  halogens,  but  it  should  be  definitely  stated  that  it 
is  never  univalent  like  the  latter. 

Occurrence,  Preparation,  and  .Properties.  —  Manganese  is  some- 
times found  uncombined  in  meteoric  iron,  otherwise  it  occurs  in 
chemical  combination  in  pyrolusite  MnO2,  braunite  Mn2O3, 
manganite  Mn2O3.H2O,  hausmannite  Mn3O4,  manganese  blende 
MnS,  and  rhodochrosite,  or  manganese  spar,  MnCO3.  The  chief 
of  these,  ores  is  pyrolusite.  The  main  localities  are  Russia, 
Brazil,  India,  Germany,  and  the  United  States.  In  small 
amounts,  manganese  is  very  widely  distributed  in  soils  and 
rocks,  also  in  traces  in  plants  and  animals.  The  annual  output 
of  manganese  ores  is  about  700,000  tons,  most  of  which  comes 
from  the  Caucasus  region. 

Manganese  is  prepared  by  heating  its  oxides  with  carbon  in 
an  electric  furnace,  or  more  readily  by  the  Goldschmidt  process 
of  igniting  the  oxides  with  aluminum. 

Manganese  is  a  hard,  steel-gray,  brittle  metal  of  specific 
gravity  8.0.  In  outward  appearance  it  is  not  unlike  cast  iron. 
It  melts  at  about  1300°.  On  exposure  to  moist  air,  it  gradually 

495 


496  OUTLINES   OF  CHEMISTRY 

assumes  a  reddish  hue  due  to  superficial  oxidation.  From  boil- 
ing  water  it  evolves  hydrogen.  In  dilute  acids  it  readily  dis- 
solves, forming  hydrogen  and  manganous  salts.  Manganese  is 
non-magnetic.  Its  atomic  weight  is  54.93. 

The  alloys  of  manganese  with  iron  are  important  in  steel 
manufacture.  The  alloys  usually  employed  are  spiegeleisen, 
which  contains  from  10  to  20  per  cent  manganese,  and  ferro- 
manganese,  which  contains  from  20  to  80  per  cent.  Alloys  of 
manganese  and  copper  contain  about  30  per  cent  manganese; 
they  are  called  manganese  bronze,  are  very  hard,  and  possess 
great  tensile  strength. 

Though  pyrolusite  MnO2  was  known  for  a  long  time,  it  was 
not  until  1774  that  its  real  nature  was  discovered  by  Scheele. 
In  1807  Gahn  reduced  the  oxide,  and  so  isolated  the  metal. 

Oxides.  — Manganese  forms  the  following  oxides:  the  monox- 
ide MnO,  the  sesquioxide  Mn2O3,  the  protosesquioxide  Mn3O4, 
the  peroxide  MnO2,  the  trioxide  or  manganic  anhydride  MnO3, 
and  the  heptoxide  or  permanganic  anhydride  Mn2O7.  Of  these 
the  first  three  are  basic  in  character.  Manganese  dioxide  yields 
manganous  salts  on  treatment  with  acids,  and  half  of  its  oxygen 
becomes  available  to  effect  oxidations.  With  strong  bases,  it 
forms  manganites.  The  oxides,  MnO3  and  Mn2O7,  are  acidic. 
On  ignition  in  the  air,  all  oxides  of  manganese  are  finally  changed 
to  Mn3O4,  which  is  probably  Mn(MnO2)2;  and  hausmannite  is 
therefore  analogous  to  the  spinels,  though  its  crystals  are 
tetragonal. 

Manganous  oxide  MnO  is  a  green  powder  formed  by  heating 
higher  oxides  in  hydrogen,  or  by  ignition  of  manganous  carbon- 
ate out  of  contact  with  the  air.  Manganous  hydroxide  is  a 
white  precipitate  formed  by  the  addition  of  caustic  alkalies  to 
solutions  of  manganous  salts.  It  readily  turns  brown  because 
of  oxidation  to  manganic  hydroxide  Mn(OH)3;  the  latter  easily 
loses  water  and  changes  to  MnO  •  OH,  which  on  careful  ignition 
yields  a  brown  powder,  manganese  sesquioxide  Mn2O3.  This 
forms  manganous  nitrate  and  manganese  dioxide  MnO2  on 
digesting  with  nitric  acid.  By  ignition  of  the  nitrate,  the  di- 
oxide may  also  be  obtained,  though  on  being  strongly  heated  it 
loses  oxygen  and  passes  over  into  Mn3O4.  Manganese  dioxide 
is  also  produced  at  the  anode  when  manganous  salts  are  electro- 
lyzed.  It  is  a  conductor  of  electricity,  and  like  lead  peroxide 


MANGANESE  497 

it  is  frequently  used  as  an  anode  in  batteries.  With  hot  hydro- 
chloric acid,  manganese  dioxide  yields  manganous  chloride  and 
chlorine;  while  in  cold  hydrochloric  acid  a  dark  brown  solution 
is  formed.  This  probably  contains  MnCl4,  which  decomposes 
into  MnCl2  and  chlorine  on  warming.  With  lime,  manganese 
dioxide  forms  manganites  of  the  composition  CaO-MnO2  and 
CaO  -2  MnO2.  These  act  like  a  mixture  of  CaO  and  MnO2,  of 
which  fact  advantage  is  taken  in  the  Weldon  process  of  again 
using  MnCl2  liquors  for  making  chlorine.  Manganese  trioxide 
is  an  unstable,  dark  red  mass.  Manganese  heptoxide  Mn2O7  is 
a  dark,  reddish  green,  oily  liquid  obtained  by  treating  potassium 
permanganate  with  concentrated  sulphuric  acid.  The  mixture 
must  be  carefully  cooled  with  ice,  for  otherwise  violent  decom- 
position will  occur. 

Salts  of  Manganese.  —  The  stable  salts  in  which  manganese  acts 
as  a  base  are  the  manganous  compounds.  In  these  the  element 
is  bivalent.  They  are  obtained  by  dissolving  any  oxide  or  hy- 
droxide of  manganese  with  an  acid.  In  the  latter  process,  the 
higher  oxides  or  hydroxides  yield  oxygen  that  is  available  to 
effect  oxidations. 

Manganous  chloride  MnCl2  •  4  H2O  forms  pink,  deliquescent, 
monoclinic  crystals.  The  salt  is  obtained  as  a  by-product  when 
chlorine  is  prepared  by  the  action  of  hydrochloric  acid  on  man- 
ganese dioxide.  Manganous  chloride  may  be  dehydrated  by 
heating  it  in  a  current  of  hydrochloric  acid  gas.  It  is  the  only 
chloride  of  manganese  that  has  been  prepared  in  pure  form. 
The  double  salt  MnCl2  •  2  NH4C1  -  H2O  may  readily  be  made.  It 
forms  crystals  of  the  isometric  system,  which  on  heating  yield 
anhydrous  manganous  chloride,  just  as  by  heating  magnesium 
ammonium  chloride,  anhydrous  magnesium  chloride  is  obtained. 

Manganous  sulphate  MnSO4  •  7  H2O  separates  from  solutions 
at  temperatures  below  6°  in  form  of  pink  monoclinic  crystals 
that  are  isomorphous  with  other  vitriols  like  FeSO4  •  7  H2O, 
ZnSO4-  7  H2O,  etc.  Between  6°  and  20°  the  salt  crystallizes  in 
triclinic  crystals  MnSO4  •  5  H2O,  that  are  isomorphous  with  blue 
vitriol  CuSO4-5H2O;  and  above  25°  orthorhombic  crystals 
MnSO4*4  H2O  are  obtained.  Manganous  sulphate  forms  double 
salts  with  alkali  sulphates,  like  K2SO4  •  MnSO4-  6  H2O.  These 
sulphates  are  isomorphous  with  similar  salts  of  magnesium,  zinc, 
nickel,  cobalt,  and  iron. 

2K 


498  OUTLINES  OF   CHEMISTRY 

Manganous  nitrate  Mn(NO3)2  -  6  H2O,  is  a  deliquescent,  pink 
salt  which  melts  in  its  water  of  crystallization  at  about  25°. 

Manganous  carbonate  MnCO3  occurs  in  nature  in  reddish, 
hexagonal  crystals  as  manganese  spar.  It  is  also  formed  as  a 
white,  insoluble  precipitate  by  adding  soluble  carbonates  to  so- 
lutions of  manganous  salts.  In  water  charged  with  carbon 
dioxide  it  dissolves,  which  behavior  is  similar  to  that  of  calcium 
carbonate. 

Manganic  chloride  MnCl3  is  supposed  to  exist  in  the  brown 
liquid  obtained  by  dissolving  manganese  dioxide  in  cold  hydro- 
chloric acid.  It  has  never  been  isolated. 

Manganic  sulphate  Mn2(SO4)3  is  a  dark  green  powder  formed 
by  heating  manganese  dioxide  with  concentrated  sulphuric  acid. 
It  is  unstable  and  readily  passes  over  into  manganous  sulphate, 
sulphur  dioxide,  and  oxygen  on  heating.  With  alkali  sul- 
phates, manganic  sulphate  forms  double  salts  that  are  isomor- 
phous  with  alum,  like  K2SO4.Mri2(SO4)3-24  H2O. 

Manganates  and  Permanganates.  —  Manganates  are  salts  of 
manganic  acid  H2MnO4,  which  has  not  been  isolated.  By  fus- 
ing manganese  dioxide  with  caustic  potash,  potassium  manga- 
nate  K2MnO4  is  formed, 'as  a  green  mass,  thus  :  — 

3  MnO2  +  2  KOH  ==  Mn2O3  +  K2MnO4  +  H2O. 

On  treating  the  mass  with  water,  a  dark  green  solution  is  ob- 
tained, from  which,  on  evaporation,  greenish  black,  rhombic 
crystals  of  K2MnO4  are  deposited  These  are  isomorphous 
with  potassium  sulphate  and  potassium  chromate.  Manganates 
of  sodium  or  potassium  may  be  obtained  by  fusing  any  oxide  or 
salt  of  manganese  with  sodium  or  potassium  carbonate  or  hydrox- 
ide, plus  sodium  or  potassium  nitrate  or  chlorate.  The  presence 
of  the  oxidizing  agent,  or  even  the  action  of  the  oxygen  of  the 
air,  insures  the  more  complete  conversion  of  the  manganese 
compound  into  manganate.  Solutions  of  manganates  are  stable 
only  when  they  contain  an  excess  of  caustic  alkali.  On  dilution 
with  water,  manganates  suffer  decomposition,  thus  :  — 

3  K2MnO4  +  2  H20  =  2  KMnO4  +  4  KOH  +  MnO2. 

This  change  is  effected  more  readily  by  passing  carbon  dioxide 
through  the  solution  of  the  manganate :  — 

3  KaMnO4  +  2  CO2  =  2  KMnO4  +  2  K2CO8  +  MnOs. 


MANGANESE  499 

The  change  may  also  be  produced  by  other  dilute  acids  or  by 
the  addition  of  chlorine. 

The  new  salt  thus  formed  is  potassium  permanganate  KMnO4. 
Its  solutions  have  a  beautiful  purple-red  color,  and  so  in  the 
reactions  just  mentioned  the  green  solution  of  the  manganate 
gradually  changes  through  blue  and  violet  to  the  characteristic 
purple  color  of  the  permanganate.  Because  of  this  change  of 
color,  Scheele  called  potassium  manganate  chameleon  mineral. 
Sometimes  potassium  permanganate  solution  is  called  chameleon 
solution,  for  on  treating  it  with  hot,  concentrated,  caustic  alkali, 
it  again  turns  green,  because  of  the  formation  of  the  manganate, 

thus : — 

2  KMn04  +  2  KOH  =  2  K2MnO4  -1-  H2O  +  O. 

Potassium  permanganate  forms  very  dark  purple,  lustrous 
crystals  of  the  rhombic  system.  They  are  isomorphous  with 
crystals  of  potassium  perchlorate  KC1O4.  On  heating  potassium 
permanganate,  it  is  decomposed :  — 

2  KMnO4  =  MnO2  +  K2MnO4  +  Oa. 

When  treated  with  concentrated  sulphuric  acid  in  the  cold,  per- 
manganic anhydride  Mn2O7  is  formed,  as  already  mentioned, 
thus : — 

2  KMn04  +  H2S04  =  Mn2O7  4-  K2SO4  +  H2O. 

This  heptoxide  is  unstable.     It  gradually  decomposes :  — 
Mn2OT  =  2  MnO2  +  O8. 

At  somewhat  elevated  temperatures,  this  reaction  proceeds  with 
explosive  violence.  The  vapors  of  the  heptoxide  are  violet. 
Paper,  alcohol,  ether,  illuminating  gas,  and  other  combustible 
substances  burst  into  flame  when  brought  in  contact  with  per- 
manganic anhydride,  because  of  the  tremendous  oxidizing  power 
of  the  ozone  that  is  being  liberated.  By  dissolving  permanganic 
anhydride  in  water  at  0°,  a  purple  red  solution  of  permanganic 
acid  HMnO4  is  obtained.  This  may  also  be  made  by  treating 
a  solution  of  barium  permanganate  with  sulphuric  acid,  for  thus 
barium  is  precipitated  as  barium  sulphate.  Though  much  more 
stable  than  manganic  acid,  nevertheless  permanganic  acid 
HMnO4-a:H2O  gradually  decomposes,  especially  on  being 
heated  or  exposed  to  light,  thus :  — 

2  HMn04  =  2  MnO2  +  H2O  +  3  O. 


500  OUTLINES   OF   CHEMISTRY 

Uses  of  Permanganates.  —  On  account  of  their  great  oxidizing 
power,  permanganates  are  used  as  disinfectants,  and  as  oxidiz- 
ing agents  in  many  chemical  processes.  They  also  serve  in  the 
preparation  and  analysis  of  many  substances  in  the  laboratory. 
Potassium  permanganate  is  commonly  employed ;  though  the 
cheaper  sodium  permanganate  NaMnO4,  which  does  not  crystal- 
lize, is  also  made  and  sold  in  solutions  as  Condy's  disinfecting 
fluid.  Wood  alcohol  is  readily  oxidized  to  formic  aldehyde  by 
potassium  permanganate,  which  fact  is  used  to  produce  formic 
aldehyde  in  fumigating  infected  houses,  etc. 

In  alkaline  or  neutral  solutions,  potassium  permanganate 
yields  oxygen  that  is  available  for  oxidation,  with  concomitant 
formation  of  manganese  dioxide  :  — 

2  KMnO4  +  H2O  =  2  KOH  +  2  MnO2  +  3  O. 

Two  molecules  of  permanganate  thus  yield  three  available  oxy- 
gen atoms.  The  oxidation  of  wood  alcohol  to  formic  aldehyde 
would  be  expressed  thus  :  — 

2  KMn04  +  3  CH3OH  =  2  KOH  +  2  MnO2+  3  HCHO  +  2  H2O. 

In  alkaline  solutions  potassium  permanganate  serves  for  the 
destruction  of  organic  matter  in  the  analysis  of  waters,  ferti- 
lizers, etc.,  the  nitrogen  present  in  the  substances  being  simul- 
taneously liberated  as  ammonia. 

When  potassium  permanganate  is  to  be  used  in  acid  solution, 
sulphuric  acid  is  commonly  employed.  Thus,  when  reducing 
substances  are  present,  a  colorless  solution  of  manganous  sul- 
phate and  potassium  sulphate  is  produced.  In  sulphuric  acid 
solutions,  two  molecules  of  potassium  permanganate  yield  jive  atoms 
of  oxygen  that  are  available  for  oxidation:  — 

2  KMnO4  +  3  H2SO4  =  K2SO4  +  2  MnSO4  +  3  H2O  +  5  O. 

Since  the  oxidation  of  oxalic  acid  to  carbon  dioxide  and  water 
requires  one  atom  of  oxygen  per  molecule  of  oxalic  acid,  we 
have : — 

(COOH)2  +  O  =  2  CO2  +  H2O. 

Consequently  if  the  oxidation  is  effected  by  means  of  potassium 
permanganate,  we  have  :  — 

2  KMnO4  +  3  H2SO4  +  5  (COOH)2 

=  10  C02  +  K2S04  +  2  MnS04  +  8  H2O. 


MANGANESE  501 

Ferrous  sulphate  is  converted  to  ferric  sulphate,  thus  :  — 
2  FeS04  +  H2S04  +  O  =  Fe2(SO4)3  +  H2O. 

Therefore,  if  the  oxidation  is  carried  on  by  means  of  potassium 
permanganate,  we  have  :  — 

2  KMnO4  +  8  H2SO4  +  10  FeSO4 

=  5  Fe2(SO4)3  +  K2SO4  +  2  MnSO4  +  8  H2O. 

With  nitrous  acid  the  reaction  is  :  — 

2  KMn04  +  3  H2SO4  +  5  HNO2 

=  K2S04  +  2  MnS04  +  5  HNO3  +  3  H2O. 

Potassium  permanganate  and  hydrogen  peroxide  mutually  re- 
duce each  other,  thus  :  — 


This  reaction  is  commonly  used  to  determine  the  content  of  a 
solution  of  hydrogen  peroxide  by  means  of  potassium  perman- 
ganate. 

Analytical  Tests  for  Manganese.  —  Manganous  sulphide  MnS 
is  precipitated  as  a  flesh-colored,  hydrous,  amorphous  substance, 
when  ammonium  sulphide  is  added  to  an  aqueous  solution  of 
any  compound  of  manganese.  This  sulphide  readily  turns  dark 
on  exposure  to  the  air,  because  oxidation  takes  place.  In  dilute 
acid,  even  in  acetic  acid,  manganous  sulphide  is  readily  soluble. 

From  solutions  of  manganous  salts,  soluble  carbonates  precipi- 
tate manganous  carbonate,  and  caustic  alkalies  precipitate  man- 
ganous hydroxide.  The  latter  is  soluble  in  solutions  of  ammonia 
or  ammonium  salts. 

Amethyst-colored  beads  are  produced  by  manganese  com- 
pounds with  either  borax  or  microcosmic  salt.  Furthermore, 
when  fused  with  a  mixture  of  soda  and  saltpeter,  a  green 
manganate  is  obtained  which,  when  dissolved  in  water  and 
treated  with  carbon  dioxide,  yields  the  characteristic  purple 
color  of  permanganate  solutions. 


CHAPTER   XXX 

IRON,   NICKEL,   AND   COBALT 

THE  atomic  weights  of  the  elements  of  this  group  are  nearly 
alike,  being  Fe — 55.84,  Ni— 58.68,  and  Co— 58.97.  More- 
over, it  will  be  recalled  that  the  atomic  weight  of  manganese, 
54.93,  is  but  little  less  than  that  of  iron.  The  chemical  simi- 
larities between  manganese  and  iron  have  been  mentioned. 
Iron  forms  ferrous  and  ferric  compounds.  In  the  former  it  is 
bivalent,  and  in  the  latter  trivalent.  Iron  is  also  hexavalent, 
forming  ferrates,  which  are  analogous  to  manganates  and  chro- 
rrates.  The  ferrous  compounds  are  analogous  to  those  of  the 
magnesium  group,  also  to  cupric  and  manganous  compounds, 
while  the  ferric  compounds  are  analogous  to  those  of  aluminum 
and  chromium.  Nickel  forms  but  one  series  of  salts.  In  these 
the  metal  is  bivalent.  Nevertheless,  a  sesquioxide  and  a  cor- 
responding hydroxide  of  nickel  are  known,  though  salts  of 
these  are  lacking.  Cobalt,  like  iron,  forms  two  series  of  com- 
pounds, the  cobaltous,  in  which  the  metal  is  bivalent,  and  the 
cobaltic,  in  which  it  is  trivalent. 

Occurrence  of  Iron.  —  Iron  is  very  abundant  and  widely  dis- 
tributed. In  meteorites  it  is  found  uncombined,  also  in  small 
grains  in  some  of  the  crystalline  rocks.  Compounds  of  iron  are 
found  in  plant  and  animal  tissues,  particularly  in  the  chloro- 
phyll of  plants  and  the  hemoglobin  of  the  blood  of  animals. 
Iron  is  necessary  for  plant  and  animal  life,  though  its  real  func- 
tion in  the  vital  processes  is  not  understood.  All  soils  and 
rocks  contain  compounds  of  iron.  Ores  of  iron  are  found  in 
enormous  quantities.  The  most  important  of  these  are  hema- 
tite Fe2O3,  magnetite,  or  magnetic  iron  ore,  Fe3O4,  limonite 
2  Fe2O3 -f  3  H2O,  and  siderite  FeCO3.  Iron  is  also  found  in 
combination  with  sulphur,  as  in  pyrite  FeS2,  but  these  com- 
pounds are  not  used  for  making  metallic  iron.  In  rocks, 
iron  is  found  in  the  form  of  oxides  and  silicates ;  and  so  by  the 
weathering  of  rocks  iron  gets  into  the  soil  and  all  natural 
waters,  whence  it  enters  plants  and  animals. 

602 


IRON,   NICKEL,   AND   COBALT 


508 


Metallurgy  of  Iron.  —  Metallic  iron  has  been  known  to 
man  for  thousands  of  years.  The  Assyrians  used  iron  knives 
and  saws,  the  ancient  Egyptians  reduced  iron  ores  and  made 
steel,  both  Homer  and  Hesiod  mention  the  forging  of  iron  for 
weapons.  Through  the  influence  of  the  Romans  iron  came 
into  more  general  use.  Still, 
the  metal  was  costly,  because 
the  processes  of  preparing  it 
from  ores  were  relatively  diffi- 
cult to  carry  out,  imperfectly 
understood,  and  known  to  but 
few;  and  so  iron  was  rather 
slow  in  replacing  bronze  for 
use  in  weapons  and  other 
implements. 

Iron  is  made  by  reducing  its 
ores  with  carbon  at  high  tem- 
peratures. The  ores,  which 
consist  of  the  oxides,  or  car- 
bonate, do  not  require  pre- 
liminary roasting.  The  re- 
duction is  effected  in  blast 
furnaces,  a  cross  section  of  a 
modern  type  of  which  is  repre- 
sented in  Fig.  151.  These 
furnaces  are  from  80  to  100 
feet  high,  and  have  a  diameter 
of  about  20  feet  where  they 
are  widest.  They  are  lined 
inside  with  fire  brick.  The 
lower  end  of  the  furnace  is 
provided  with  openings  of 
tubes,  the  tuyeres,  through 
which  hot  air  can  be  forced 
into  the  furnace.  The  latter  is  heated  and  then  charged  from 
the  top  with  ore  properly  mixed  with  coke  and  limestone. 
Instead  of  coke  anthracite  coal  or  charcoal  may  be  used.  The 
purpose  of  the  limestone  is  to  form  calcium  silicates,  or  so-called 
slag,  with  the  sand  present  in  the  ore.  If  the  latter  contains 
carbonates  of  calcium  and  magnesium,  or  other  basic  materials, 


FIG.  151. 


504  OUTLINES   OF   CHEMISTRY 

sand  is  added  instead  of  limestone  to  make  the  slag.  In  any 
case  the  material  added  to  the  ore  to  produce  the  fusible  slag 
is  termed  the  flux.  The  charge  is  carried  to  the  top  of  the 
furnace  by  some  form  of  mechanical  conveyor,  and  is  introduced 
through  a  bell  trap,  so  arranged  that  while  the  material  is  put 
into  the  furnace  practically  no  gases  escape  from  the  latter. 
As  air  heated  to  about  8009  is  blown  up  through  the  charge, 
the  latter  becomes  very  hot.  In  the  lower  part  of  the  furnace 
carbon  dioxide  forms,  due  to  the  combustion  of  the  coke. 
The  carbon  dioxide  as  it  rises  through  the  hot  layers  of  coke  is 
reduced  to  carbon  monoxide,  and  the  latter  acts  on  the  ferric  oxide 
and  reduces  it  to  iron,  thus :  — 

Fe203  +  3  CO  =  2  Fe  +  3  CO2. 

The  gases,  still  rich  in  carbon  monoxide,  pass  out  through  the 
vent  at  the  top  of  the  furnace,  and  are  used  as  fuel  to  heat  the 
air  before  it  is  blown  into  the  furnace.  Of  late  they  are  also 
frequently  used  to  run  gas  engines.  The  slag  and  iron  settle 
to  the  lower  part  of  the  furnace,  forming  two  layers,  the  liquid 
slag  covering  the  heavier  molten  iron.  The  iron  is  tapped  off 
from  below  and  run  into  molds  of  sand  about  every  eight  hours. 
In  this  way  rough  bars  of  cast  iron  called  pigs  are  produced. 
The  slag  runs  continually  from  a  lateral  opening  above  the 
iron.  Blast  furnace  slag  is  now  frequently  used  for  making 
Portland  cement.  As  the  material  in  the  furnace  melts  down, 
the  latter  is  fed  from  the  top,  so  that  the  process  once  started 
is  continuous,  furnaces  remaining  in  operation  for  years  at  a 
time.  In  the  upper  parts  of  the  furnace  the  temperature  is 
not  high  enough  to  melt  the  iron,  which  is  formed  in  a  porous 
or  spongy  condition,  and  carried  down  with  the  slag  to  the 
lower  and  much  hotter  parts.  Here  it  takes  up  carbon,  form- 
ing iron  carbide  in  part,  which  with  the  iron  yields  a  mixture 
that  melts  at  a  much  lower  temperature  than  iron  free  from 
carbon.  It  is  this  iron  containing  carbon  and  other  impurities 
that  is  tapped  from  the  furnace  and  made  into  pig  iron.  The 
latter  is  essentially  the  same  as  cast  iron,  besides  which  we 
also  have  wrought  iron  and  steel,  as  the  main  varieties  of  iron. 
All  kinds  of  iron  used  in  practice  contain  carbon  and  various 
other  impurities,  pure  iron  being  practically  unknown. 

Cast  Iron.  —  The  melting  point  of  cast  iron  varies  from  about 


IRON,    NICKEL,   AND   COBALT  505 

1050°  to  1300°,  according  to  its  content  of  carbon  and  other 
impurities.  Cast  iron  contains  from  2.3  to  5  per  cent  carbon, 
besides  smaller  amounts  of  silicon,  phosphorus,  sulphur,  and 
manganese.  When  cooled  very  slowly,  most  of  the  carbon  in 
cast  iron  crystallizes  out  in  the  form  of  leaflets  of  graphite. 
This  iron  consequently  appears  dark  gray  on  its  fractured  surface 
and  is  known  as  gray  cast  iron.  It  is  used  in  making  castings, 
for  it  has  a  low  melting  point,  contracts  uniformly  on  cooling, 
and  can  afterwards  readily  be  worked  with  tools.  On  dissolv- 
ing gray  cast  iron  in  hydrochloric  acid,  the  graphitic  carbon 
remains  behind,  while  the  carbon  which  was  combined  with  the 
iron  in  the  form  of  carbides  is  given  off  with  the  hydrogen,  being 
evolved  as  hydrocarbon  gases.  Good  gray  castings  contain 
about  2  to  3  per  cent  of  graphitic  carbon,  and  1  to  1.5  per  cent 
combined  carbon. 

On  cooling  cast  iron  rapidly,  practically  all  the  carbon  re- 
mains combined  with  the  iron,  forming  carbides.  Such  iron 
has  a  silver-white  fracture  and  is  known  as  white  cast  iron.  It 
is  very  hard  and  brittle,  and  is  consequently  not  employed  for 
castings,  but  is  converted  into  wrought  iron.  Cast  iron  con- 
taining from  5  to  20  per  cent  manganese  takes  up  considerable 
amounts  of  carbon.  Its  fracture  shows  a  coarsely  crystalline 
structure,  whence  the  iron  is  known  as  spiegeleisen.  Ordi- 
nary cast  iron  contains  from  0.5  to  4  per  cent  silicon,  from  0.4 
to  2  per  cent  phosphorus,  and  sulphur  varying  up  to  about  0.2 
per  cent.  Sulphur  and  phosphorus  make  iron  brittle  and  hence 
are  quite  objectionable. 

Wrought  Iron.  —  This  is  nearly  pure  iron,  often  containing 
less  than  1  per  cent  of  impurities.  It  is  made  by  puddling 
cast  iron.  This  process  consists  of  heating  the  pig  iron  with 
iron  oxide  in  a  current  of  air  on  the  hearth  of  a  reverberatory 
furnace.  Thus  the  impurities  in  the  iron  are  oxidized.  The 
carbon  largely  escapes  as  carbon  monoxide.  The  silicon  and 
phosphorus  after  oxidation  unite  with  some  of  the  iron,  yielding 
a  slag,  which  also  contains  the  sulphur. 

As  the  heating  proceeds,  the  molten  iron  becomes  more  and 
more  viscous.  It  is  continually  stirred  or  puddled  so  that  the 
air  may  gain  access  to  it.  Finally,  the  mass  becomes  so  thick 
that  it  can  be  worked  up  into  a  ball,  which  is  then  taken  from 
the  furnace  and  rolled,  or  hammered  by  a  steam  hammer. 


506  OUTLINES   OF   CHEMISTKY 

Thus  the  slag  is  removed  from  the  iron,  and  a  malleable,  duc- 
tile product  is  obtained  which  often  contains  less  than  0.2  per 
cent  carbon.  Wrought  iron  melts  at  about  1600°,  and  is  plastic 
enough  to  be  welded  at  from  900°  to  1100°.  In  welding,  the 
two  ends  to  be  joined  are  brought  to  the  welding  temperature 
in  a  forge.  Borax  or  sand  is  sprinkled  over  the  parts,  which 
are  then  again  heated  for  a  few  moments  till  the  borax  or  sand 
has  formed  a  slag  with  the  oxides  of  iron  on  the  surface.  This 
slag  is  a  borate  or  silicate  of  iron.  It  protects  the  iron  from 
oxidation.  On  now  hammering  the  parts  together,  the  slag 
flies  off,  and  the  iron  coming  into  actual  contact  welds.  Be- 
cause of  its  very  low  carbon  content,  wrought  iron  will  not  harden 
when  rapidly  chilled,  as  does  cast  iron  or  steel. 

By  heating  iron  castings  covered  with  pulverized  iron  ore 
for  about  48  hours,  and  allowing  them  to  cool  slowly,  they 
become  sufficiently  malleable  for  many  purposes.  The  process 
abstracts  some  of  the  carbon  from  the  iron,  which  then  is  called 
malleable  iron.  It  is  much  cheaper  than  wrought  iron,  in  place 
of  which  it  is  often  used  when  possible. 

Steel. — Steel  contains  more  carbon  than  wrought  iron,  but 
much  less  than  cast  iron.  It  also  contains  practically  no  sul- 
phur or  phosphorus,  and  ordinarily  runs  low  in  silicon.  The 
amount  of  carbon  in  steel  varies  from  about  0.2  to  1.6  per  cent. 
That  which  contains  the  least  carbon,  0.2  per  cent,  approaches 
wrought  iron  in  quality  and  is  called  mild  steel.  Steel  used 
for  building  purposes  is  termed  structural  steel.  It  contains 
from  0.2  to  0.8  per  cent  carbon,  whereas  tool  steel  has  a  higher 
carbon  content,  namely  from  about  0.8  to  1.5  per  cent.  Like 
cast  iron,  steel  may  be  hardened  by  heating  and  then  suddenly 
cooling  it.  On  the  other  hand,  when  such  hardened  steel  is 
again  heated  to  redness  and  cooled  slowly,  the  material  is  soft. 
As  in  the  case  of  cast  iron,  the  heating  and  sudden  chilling 
leaves  the  carbon  in  the  combined  condition  and  thus  makes  a 
hard  steel;  whereas  on  slow  cooling,  the  carbon  crystallizes  out 
as  graphite  and  so  produces  a  soft,  pliable  product.  By  proper 
heating  and  cooling,  steel  of  any  desired  hardness  may  be  pro- 
duced. This  process  is  called  tempering. 

Pig  iron  is  converted  into  steel  by  either  the  Bessemer  process 
or  the  open  hearth  or  Siemens- Martin  process.  In  the  Bessemer 
process  the  molten  cast  iron  is  poured  into  a  converter  (Fig.  152), 


IRON,   NICKEL,    AND   COBALT 


507 


through  whose  perforated  bottom  compressed  air  is  then  blown 
into  the  metal ;  thus  carbon,  silicon,  and  phosphorus  are  oxi- 
dized. Molten  spiegeleisen  of  known  carbon  content  is  then 
added  in  proper  quantity  to  produce  steel  of  the  exact  carbon 
content  desired.  The  converter  is  mounted  so  that  it  can  be 
tilted  to  pour  out  or  receive  its  contents.  The  slag,  which  con- 
sists of  silicates  of  manganese,  iron,  calcium,  and  sulphide  of 
iron,  is  first  poured  off,  after  which  the  molten  steel  is  run  into 
molds  to  form  ingots.  These  are  afterwards  rolled  into  rails. 


FIG.  152. 


By  means  of  a  Bessemer  converter  about  20  tons  of  cast  iron 
can  be  converted  into  steel  in  20  to  30  minutes.  The  converter 
is  made  of  iron  and  is  usually  lined  with  material  similar  to 
that  of  fire  brick.  In  case  the  cast  iron  is  rich  in  phosphorus, 
this  siliceous  lining  is  replaced  by  one  of  calcium  and  magnesium 
oxides  obtained  by  calcining  dolomite.  This  basic  lining  absorbs 
the  phosphorus,  yielding  calcium  and  magnesium  phosphates, 
which  appear  in  the  slag.  This  adaptation  of  the  Bessemer 
process  to  the  treatment  of  cast  iron  rich  in  phosphorus  is 
know  as  the  Thomas-Gilchrist  process,  and  the  slag  produced 
by  it  is  called  the  Thomas  slag.  TJie  latter,  being  rich  in  phos- 
phates, is  ground  up  and  sold  as  a  fertilizer.  Thus,  ores  rich  in 


508  OUTLINES  OF  CHEMISTRY 

phosphorus,  which  were  formerly  discarded  as  unfit  for  Bes- 
semer steel,  are  very  profitably  turned  to  use. 

The  open  hearth  or  Siemens-Martin  process  consists  essentially 
of  heating  the  cast  iron  together  with  rusty  scrap  iron  or  other 
iron  oxide,  commonly  hematite  ore,  on  the  hearth  of  a  special 
type  of  reverberatory  furnace,  using  gas  as  fuel.  The  oxida- 
tion of  the  carbon  in  the  cast  iron  thus  proceeds  at  the  expense 
of  the  oxygen  in  the  iron  oxide.  The  process  is  continued  till 
a  sample  taken  from  the  material  shows  that  the  oxidation  has 
proceeded  far  enough  to  produce  the  steel  desired.  The 
molten  mass  is  then  run  off  into  molds.  The  process  requires 
about  8  hours  for  its  completion;  but  it  yields  an  excellent, 
uniform  steel,  and  utilizes  the  carbon  in  the  cast  iron  for  the 
reduction  of  ores.  Hence  it  is  more  economical  than  the  Bes- 
semer process,  which  it  is  rapidly  displacing.  Moreover,  if  the 
oast  iron  is  rich  in  phosphorus,  a  hearth  lined  with  the  oxides  of 
magnesium  and  calcium  may  be  used  to  absorb  the  phosphorus  in 
form  of  phosphates;  with  this  modification  the  process  is  called 
the  basic  open  hearth  process. 

In  1910  the  United  States  produced  27,303,567  tons  of  pig 
iron,  a  large  portion  of  which  was  converted  into  steel.  Nickel, 
manganese,  chromium,  silicon,  tungsten,  and  molybdenum  are 
often  added  to  steel  in  small  proportions,  thus  forming  alloys 
of  mechanical  properties  that  are  desirable  for  certain  special 
purposes. 

Properties  of  Iron. — Pure  iron  may  be  made  by  strongly  ig- 
niting ferric  oxide  in  a  current  of  hydrogen.  If  heated  thus 
to  not  higher  than  450°,  the  iron  obtained  contains  hydrogen 
and  is  called  pyrophoric  iron,  for  it  burns  spontaneously  on  ex- 
posure to  the  air.  Electrolytic  iron,  obtained  by  electrolyzing  a 
ferrous  sulphate  solution,  using  a  thick  wrought  iron  anode  and 
a  thin  iron  foil  as  cathode,  is  also  nearly  pure;  but  it  contains 
a  few  hundredths  of  one  per  cent  of  hydrogen,  which  renders  it 
very  hard  and  brittle.  The  melting  point  of  pure  iron  has  not 
been  determined  with  certainty;  it  lies  above  that  of  wrought 
iron.  The  specific  gravity  of  iron  is  7.86.  The  metal  is  white, 
malleable,  ductile,  and  fairly  soft.  It  is  attracted  by  a  magnet, 
and  becomes  magnetic;  but  it  loses  its  magnetism  rapidly, 
while  magnetized  steel  does  not.  Iron  remains  unchanged  in 
dry  air,  while  in  moist  air  or  in  presence  of  air  and  salt  solutions 


IRON,    NICKEL,   AND   COBALT  509 

it  rusts,  forming  hydrated  ferric  oxide.  This  corrosion  is 
hastened  by  local  electrolytic  action,  and  is  generally  guarded 
against  by  painting  or  tarring  the  exposed  parts  of  the  iron  or 
steel.  Dilute  hydrochloric  or  sulphuric  acid  dissolves  iron, 
forming  ferrous  chloride  or  sulphate  and  hydrogen,  which  in 
case  of  cast  iron  or  steel  is  mixed  with  hydrocarbon  gases  and 
compounds  of  carbon,  hydrogen,  phosphorus,  and  sulphur  that 
have  a  bad  odor. 

When  dipped  into  very  concentrated  nitric  acid  and  then 
rinsed,  iron  no  longer  dissolves  in  nitric  acid,  nor  does  it  pre- 
cipitate copper  from  solutions  of  its  salts.  The  iron  is  said  to 
be  passive.  It  is  thought  by  many  that  the  phenomenon  is  due 
to  a  very  thin,  invisible  coating  of  oxide  on  the  iron,  for  when 
passive  iron  is  scratched  with  a  hard  point,  it  at  once  dissolves 
rapidly  in  nitric  acid.  Other  metals,  like  chromium  and  nickel, 
also  exhibit  the  phenomena  of  the  passive  state. 

Oxides  and  Hydroxides  of  Iron. —  Ferrous  oxide  FeO  is  ob- 
tained as  a  black  powder  by  heating  ferric  oxide  in  hydrogen 
or  carbon  monoxide  to  about  300°.  On  exposure  to  the  air  it 
readily  oxidizes  further. 

Ferrous  hydroxide  Fe(OH)2  is  a  white  precipitate  formed  by 
adding  caustic  alkali  to  a  solution  of  a  ferrous  salt.  It  readily 
oxidizes  in  the  air,  turning  green  and  then  brown. 

Ferric  oxide  Fe2O3  is  found  in  nature  as  hematite  ore.  It  is 
found  in  large  quantities  in  the  Lake  Superior  region,  and  is  the 
most  important  of  the  iron  ores.  It  crystallizes  in  very  dark  red, 
hexagonal  pyramids  and  prisms.  When  finely  ground,  ferric 
oxide  is  used  as  a  pigment  in  paints  under  the  name  of  red  ocher 
or  Venetian  red.  By  ignition  of  ferrous  sulphate  or  oxalate  in 
the  air,  ferric  oxide  is  obtained,  which  when  finely  ground  is 
sold  as  a  pigment  or  as  rouge  for  polishing  purposes.  Ferric 
oxide  is  prepared  in  the  laboratory  by  ignition  of  ferric  hy- 
droxide, which  is  obtained  by  precipitating  a  solution  of  a  ferric 
salt  with  caustic  alkali. 

Ferric  hydroxide  Fe(OH)3  is  an  amorphous,  brown,  flocculent 
precipitate  to  which  alkali  adheres  very  tenaciously.  It  dis- 
solves to  some  extent  in  a  concentrated  solution  of  ferric 
chloride,  forming  a  very  dark  brown  solution  of  a  basic  ferric 
chloride;  from  this  the  chlorine  may  be  removed  in  the  form  of 
ferric  chloride  by  dialysis,  thus  leaving  a  dark  brown,  tasteless, 


510  OUTLINES   OF   CHEMISTRY 

colloidal  solution  of  ferric  hydroxide  or  so-called  dialyzed  iron 
behind. 

Hydrated  ferric  oxides  occur  in  nature.  Thus  we  have  li- 
monite  2  Fe2O3  +  3  H2O  or  Fe2O3  •  2  Fe(OH)3,  also  known  as 
brown  iron  ore;  pyrosiderite  FeO  •  OH,  and  log  iron  ore 
Fe2O(OH)4.  These  may  all  be  considered  as  dehydration 
products  of  Fe(OH)3.  Yellow  ocher  is  an  impure  hydrated 
ferric  oxide.  It  is  used  as  a  pigment  in  paints.  Though  the 
iron  oxide  pigments  are  not  as  bright  in  color  as  many  others, 
yet  they  are  valuable  because  they  are  permanent  and  cheap. 

Ferrous  ferric  oxide  Fe3O4  also  known  as  magnetic  iron  oxide, 
is  formed  as  the  final  product  of  continuous  strong  ignition  of  any 
oxide  of  iron  in  the  air.  In  nature  it  occurs  in  black  octahedra 
and  dodecahedra  as  magnetite  ore.  It  is  often  magnetic  and 
is  then  called  lodestone.  The  hammer  black  formed  as  a  scale 
on  iron  when  it  is  heated  in  the  air  is  Fe3O4.  Ferrous  ferric 
oxide  is  isomorphous  with  spinel,  and  consequently  probably 
is  Fe(FeO2)2. 

When  iron  turnings  are  fused  with  potassium  nitrate,  or 
when  chlorine  is  conducted  through  a  cold,  concentrated  solu- 
tion of  caustic  potash  containing  ferric  hydroxide  in  suspen- 
sion, potassium  ferrate  K2FeO4  is  formed.  It  may  be  obtained 
in  the  form  of  dark  red  crystals,  which  are,  however,  unstable. 
They  decompose,  yielding  oxygen,  ferric  hydroxide,  and  caustic 
potash.  Barium  ferrate  BaFeO4  is  more  stable.  The  ferrates 
are  salts  of  ferric  acid  H2FeO4,  which,  like  manganic  acid,  is 
not  known  in  the  free  state.  Ferrates  are  analogous  to  and 
isomorphous  with  sulphates  and  chromates.  When  caustic  alkali 
solutions  are  electrolyzed,  ferrates  are  formed  at  the  anode  if 
the  latter  consists  of  iron. 

Chlorides  of  Iron. — Ferrous  chloride  FeCl2  is  obtained  as  a 
white  mass  by  heating  iron  filings  in  a  current  of  hydrochloric 
acid  gas.  It  may  also  be  obtained  in  the  form  of  green,  mono- 
clinic  crystals  FeCl2-4H2O,  from  aqueous  solutions  carefully 
kept  from  the  oxygen  of  the  air,  for  the  salt  is  readily  oxidized,, 

thus :  — 

6  FeCl2  +  30  =  4  FeCl3  +  Fe2O3. 

Ferric  chloride  FeCl3  is  obtained  in  the  form  of  very  dark 
green,  lustrous,  hexagonal  crystals  when  iron  is  heated  in  a 
current  of  chlorine.  The  product  may  be  sublimed.  It  is  very 


IRON,   NICKEL,   AND   COBALT 


511 


deliquescent.  Ferric  chloride  may  readily  be  formed  by  pass- 
ing chlorine  into  a  solution  of  ferrous  chloride,  by  boiling  the 
latter  with  aqua  regia,  or  by  dissolving  ferric  oxide  or  hydrox- 
ide in  hydrochloric  acid.  On  evaporation,  a  dark  brown  crys- 
talline mass,  Fe2Cl6  •  12  H2O,  is  obtained.  At  higher  tempera- 
tures different  hydrates  crystallize  out  and  are  in  equilibrium 
with  the  saturated  solution.  This  question  has  been  carefully 
studied  by  H.  W.  Bakhuis  Roozeboom,  whose  results  are  shown 


500 


& 

cc 

SQ 

I 

I 

J 

00 

rt 
I 


-t* 


oa 
g 


ro 
o 


ntra 

— 
§ 


Fc 


:C167H 


Fe, 


Cl. 


5H2 


4H2q 


60        40 


20         0        20 

Temperature  in  degrees  C. 
FIG.  153. 


40        60 


80 


in  Fig.  153,  which  gives  the  solubility  curves  of  ferric  chlo- 
ride and  indicates  the  ranges  of  temperature  at  which  the  dif- 
ferent hydrates  are  in  equilibrium  with  the  solutions  (compare 
the  case  of  magnesium  chloride,  Fig.  136).  Ferric  chloride 
is  used  in  medicine.  The  salt  is  soluble  in  alcohol,  ether,  and 
many  other  liquids  besides  water. 

Ferrous  bromide  FeBr2  and  ferric  bromide  FeBr3  are  analo- 
gous to  the  corresponding  chlorides. 

Ferrous  iodide  FeI2  •  4  H2O  consists  of  bluish  green,  mono- 
clinic  crystals  formed  by  heating  iron  filings  and  iodine  to- 
gether under  water.  The  salt  is  used  in  medicine  in  sirup  of 
ferrous  iodide.  Ferric  iodide  is  not  known. 


512  OUTLINES   OF   CHEMISTRY 

Sulphides  and  Sulphates  of  Iron.  —  Ferrous  sulphide  FeS  is 
formed  by  heating  iron  and  sulphur  together,  or.  by  adding 
ammonium  sulphide  to  a  solution  of  a  salt  of  iron.  If  a  ferric 
salt  is  used,  a  mixture  of  sulphur  and  ferrous  sulphide  is 
obtained,  thus :  — 

2  FeCl3  +  3  (NH4)2S  =  6  NH4C1  +  2  FeS  +  S. 

Ferrous  sulphide  is  also  formed  by  warming  finely  divided  iron 
with  sulphur  in  water.  The  fused  sulphide  forms  a  black, 
brittle,  crystalline  mass  on  cooling;  whereas  the  precipitated 
sulphide  is  black  and  amorphous.  Ferrous  sulphide  is  soluble 
in  dilute  acids,  hence  it  is  not  precipitated  by  hydrogen  sulphide 
from  acid  solutions  of  salts  of  iron. 

Ferric  sulphide  Fe2S3  is  a  greenish  yellow  mass,  obtained 
by  fusing  iron  and  sulphur,  or  ferrous  sulphide  and  sulphur, 
together  in  proper  proportions.  It  is  not  formed  by  precipi- 
tating ferric  salts  with  ammonium  sulphide ;  for  the  latter  first 
reduces  the  ferric  salt,  and  then  precipitates  ferrous  sulphide. 

Iron  disulphide  FeS2  occurs  in  nature  as  pyrite  and  marcasite. 
It  may  be  made  artificially  by  carefully  heating  iron  and  sul- 
phur together  in  proper  proportions. 
Pyrite  forms  golden  yellow  crystals  hav- 
ing a  metallic  luster.  As  pyrite,  these 
crystals  are  cubes,  octahedra,  or  pentago- 
nal dodecahedra  (Fig,  154) ;  and  as  mar- 
casite, they  are  orthorhombic. 

All  sulphides  of  iron  when  roasted  in 
the  air  finally  yield  sulphur  dioxide  and 
ferrous  ferric  oxide  Fe3O4.     On  exposure 
to  moist  air  the  sulphides  are  gradually  oxidized  to  sulphates. 

Ferrous  sulphate  FeSO4-7H2O,  also  called  green  vitriol  or 
copperas,  is  formed  by  gently  roasting  pyrite  so  as  to  form 
ferrous  sulphide,  and  then  allowing  the  latter  to  oxidize  to 
sulphate  in  moist  air.  From  the  mass  the  salt  is  readily 
extracted  with  water,  and  from  the  solutions  green,  monoclinic 
prisms,  FeSO4-7H2O  (Fig.  74),  are  obtained.  These  are 
isomorphous  with  other  vitriols.  The  salt  FeSO4  •  5  H2O  is 
also  known.  It  forms  triclinic  crystals  that  are  isomorphous 
with  CuSO4-5H2O  (Fig.  75). 

Ferrous  sulphate  is  used  as  a  reducing  agent,  as  a  mordant, 


IRON,   NICKEL,   AND   COBALT  513 

as  a  disinfectant,  and  also  in  making  ordinary  writing-ink.  The 
latter  consists  essentially  of  a  solution  of  ferrous  sulphate  and 
extract  of  nutgalls,  which  contains  tannin.  Thus  a  ferrous  tan- 
nate  is  formed,  which  on  exposure  to  the  air  is  oxidized  to  a 
ferric  compound  that  is  black  and  not  readily  soluble.  Dextrine 
or  gum  arabic  is  generally  added  to  ink  to  retard  the  precipita- 
tion of  the  ferric  tannate  as  the  ink  stands  in  bottles.  An 
antiseptic  like  carbolic  acid  or  corrosive  sublimate  is  often  also 
introduced  to  prevent  the  growth  of  molds. 

Ferrous  sulphate  forms  double  salts  with  alkali  sulphates. 
Of  these  ferrous  ammonium  sulphate  (NH4)2SO4-FeSO4-6H2O 
is  of  special  importance.  It  is  known  as  Mohr's  salt  and  is  used 
in  chemical  analysis.  On  exposure  to  the  air  this  salt  does  not 
oxidize  as  readily  as  ferrous  sulphate. 

Ferric  sulphate  Fe2(SO4)3  is  formed  by  the  oxidation  of  fer- 
rous sulphate,  or  by  dissolving  ferric  hydroxide  in  sulphuric 
acid.  In  an  anhydrous  condition  it  is  a  white  mass.  This 
dissolves  in  water,  slowly  yielding  a  brown  solution.  With 
alkali  sulphates  it  forms  ferric  alums;  examples  of  these  are, 
(NH4)2S04.Fe2(S04)3 .  24H2O  and  K2SO4  •  Fe2(SO4)3-  24  H2O. 

Ferrous  Carbonate  FeCO3  occurs  in  rhombohedral  crystals  in 
nature  as  siderite.  It  is  formed  as  a  white  precipitate  when 
alkali  carbonates  are  added  to  solutions  of  ferrous  salts.  Fer- 
rous carbonate,  like  calcium  carbonate,  is  soluble  in  water 
charged  with  carbon  dioxide,  thus  forming  ferrous  bicarbonate 
Fe(HCO3)2.  On  exposure  to  the  air  ferrous  carbonate  soon 
turns  dark  in  color  because  of  the  formation  of  hydrated  ferric 
oxide.  Ferric  carbonate  is  unknown.  When  treated  with  sodium 
carbonate,  ferric  salts  yield  ferric  hydroxide.  The  reaction  is 
similar  to  that  of  aluminum  salts,  thus :  — 

Fe2(S04)3+3Na2C03+3H20  =  3Na2S04+2Fe(OH)3+3C02. 
Al2(S04)3-|-3Na2C03+3H20  =  3Na2S04  +  2Al(OH)3  +  3C02. 

Cyanides  of  Iron.  —  The  simple  cyanides  Fe(CN)2  and 
Fe(CN)3  are  unknown,  but  many  double  cyanides  of  iron 
have  been  prepared.  Potassium  ferrocyanide  K4Fe(CN)6,  or 
4KCN-Fe(CN)2,  is  made  by  fusing  together  potash,  scrap 
iron,  and  animal  refuse,  like  blood,  hoofs,  horns,  scraps  of 
hides,  etc.,  as  stated  in  Chapter  XIV.  On  cooling  and  leach- 
ing out  the  cake  with  water,  a  yellow  solution  is  obtained 

2L 


514  OUTLINES  OF  CHEMISTRY 

which  deposits  beautiful  lemon-yellow,  monoclinic  crystals 
K4Fe(CN)6  •  3  H2O.  These  are  also  known  as  yellow  prussiate 
of  potash.  They  readily  lose  water,  and  on  further  heating 
they  decompose,  yielding  potassium  cyanide,  as  already  stated. 
Potassium  ferricyanide  K8Fe(CN)6,  or  3  KCN  •  Fe(CN)3,  is 
formed  by  treating  potassium  ferrocyanide  with  chlorine,  thus :  — 

2  K4Fe(CN)6  +  C12  =  2  KC1  +  2  K3Fe(CN)6. 

The  salt  consists  of  dark  red,  rhombic  prisms  which  readily 
dissolve  in  about  three  parts  of  water,  yielding  a  greenish  brown 
solution.  The  compound  is  also  known  as  red  prussiate  of 
potash. 

When  treated  with  concentrated  hydrochloric  acid,  a  satu- 
rated solution  of  potassium  ferrocyanide  yields  a  white  crystal- 
line precipitate  which  is  H4Fe(CN)6,  i.e.  the  free  ferrocyanic 
acid.  Similarly,  ferricyanic  acid  H3Fe(CN)6  may  be  obtained 
from  K3Fe(CN)6. 

While  the  alkali  salts  of  these  acids  are  soluble,  the  salts  they 
form  with  the  heavy  metals  are  insoluble,  and  may  consequently 
be  obtained  by  precipitation  from  solutions.  Copper  ferrocy- 
anide, for  instance,  is  an  insoluble  ferrocyanide.  It  has  already 
been  described.  When  a  ferric  salt  is  added  to  a  solution  of 
potassium  ferrocyanide,  ferric  ferrocyanide  Fe'"4  [Fe"(CN)6]8  is 
formed  as  an  indigo-blue  precipitate  called  Prussian  blue:  — 

4  FeCl8  +  3  K4Fe(CN)6  =  12  KOI  +  Fe%  [Fe"(CN)6]8. 

If  the  ferric  salt  is  added  to  a  large  excess  of  potassium  ferro- 
cyanide, soluble  Prussian  blue  K2Fe'"2  [Fe"(CN)6]2  is  formed. 
Caustic  alkali  decomposes  Prussian  blue  :  — 

Fe'f'4  [Fe"(CN)6]8  +  12  KOH  =  3  K4Fe(CN)6  +  4  Fe(OH)8. 

Prussian  blue  is  used  as  a  pigment. 

When  a  ferrous  salt  is  added  to  a  solution  of  potassium  fer- 
ricyanide, ferrous  ferric  cyanide  Fe"8  [Fe'"(CN)6]a  is  precipi- 
tated: — 

3  FeS04+  2  K3Fe(CN)6=  3  K2SO4  +  Fe"3[Fe'"(CN)6]2. 

This  precipitate  is  also  indigo-blue  in  color.  It  is  known  as 
Turnbull's  blue.  When  treated  with  caustic  alkali,  it  is  decom- 
posed into  ferrous  hydroxide  and  potassium  ferricyanide. 


IRON,   NICKEL,   AND   COBALT  515 

Blue  Printing.  —  When  a  ferric  salt  i's  added  to  a  solution  of 
potassium  ferricyanide,  a  brown  solution,  but  no  precipitate,  is 
formed.  Paper  treated  with  such  a  solution  and  dried  in  the 
dark  is  the  sensitive  paper  used  in  blue  printing.  On  exposing 
this  paper  to  the  light,  the  ferric  salt  is  partially  reduced  to  the 
ferrous  state.  When  the  paper  is  then  washed  with  water,  the 
insoluble  Turnbull's  blue  formed  remains,  while  the  places  pro- 
tected from  the  light  appear  white,  because  the  original  mix- 
ture on  the  paper  is  simply  dissolved  away.  As  a  rule,  ferric 
ammonium  citrate  and  potassium  ferricyanide  are  used  in  mak- 
ing blue  print  paper.  Ammonium  hydroxide  or  caustic  alkalies 
dissolve  both  Turnbull's  and  Prussian  blue,  and  these  serve  as 
inks  to  write  white  characters  on  blue  prints. 

Other  Compounds  of  Iron.  —  Ferrous  nitrate  Fe(NO3)2  and  fer- 
ric nitrate  Fe(NO3)3  have  also  been  prepared.  The  latter  is 
deliquescent  and  yields  basic  nitrates  on  boiling  with  water. 

Ferrous  phosphate  Fe3(PO4)2  is  an  insoluble  white  precipi- 
tate, while  ferric  phosphate  FePO4  is  a  yellowish  white  powder, 
insoluble  in  water  and  acetic  acid. 

Iron  carbide  Fe3C,  also  called  cementite,  occurs  in  cast  iron 
and  steel  as  so-called  combined  carbon. 

The  silicides  Fe2Si  and  FeSi  have  been  obtained  in  crystalline 
form.  They  are  hard  and  brittle. 

The  phosphides  Fe3P  and  Fe2P  are  known.  Their  presence 
in  iron  makes  it  very  brittle. 

Ferric  acetate  Fe(C2H3O2)3  is  unstable  and  readily  hydro- 
lyzed,  forming  acetic  acid  and  basic  ferric  acetates,  which,  being 
insoluble  in  acetic  acid,  enable  the  analytical  chemist  to  pre- 
cipitate iron  from  an  acid  solution.  This  is  of  consequence  in 
separating  iron  from  manganese,  for  instance. 

Analytical  Tests  for  Iron.  —  From  solutions  of  both  ferrous 
and  ferric  salts,  ammonium  sulphide  precipitates  black  ferrous 
sulphide,  which  readily  dissolves  in  acids. 

In  solutions  of  ferrous  salts,  which  are  almost  colorless, 
hydroxides,  carbonates,  and  ferricyanides  of  the  alkalies  pro- 
duce characteristic  precipitates  that  have  already  been  de- 
scribed. In  solutions  of  ferric  salts,  which  are  commonly 
brown,  hydroxides,  carbonates,  and  ferrocyanides  of  the  alkalies 
also  produce  characteristic  precipitates;  while  potassium  sul- 
phocyanate  KCNS  forms  a  deep  red  coloration  of  soluble  ferric 


516  OUTLINES  OF   CHEMISTRY 

sulphocyanide  Fe(CNS)3.  In  the  borax  bead,  ferrous  com- 
pounds yield  a  green,  and  ferric  compounds  a  brown,  coloration. 

The  fact  that  ferric  salts  may  readily  be  changed  into  ferrous 
salts  by  many  reducing  agents,  like  nascent  hydrogen,  stannous 
chloride,  hydrogen  sulphide,  etc.,  is  frequently  used  in  analyt- 
ical chemistry.  The  change  of  ferrous  salts  to  the  ferric  state 
by  oxidizing  agents  like  nitric  acid,  bichromates,  or  perman- 
ganates, is  also  often  employed. 

Occurrence,  Preparation,  and  Properties  of  Nickel.  — Nickel  is 
found  in  meteoric  iron.  The  chief  ores  of  nickel  are  nicollite 
NiAs,  gersdorffite,  or  nickel  glance,  NiAsS,  and  garnierite 
Mg2Ni2H4(SiO4)3  .  4  H2O.  The  chief  localities  are  Ontario 
and  New  Caledonia. 

The  commercial  production  of  nickel  is  a  complicated  pro- 
cess which  will  not  be  described  here.  Pure  nickel  may  be 
obtained  by  igniting  the  oxides  or  the  oxalate  in  a  current  of 
hydrogen  ;  also  by  Goldschmidt's  process,  or  by  reducing  the 
oxides  with  carbon. 

Nickel  is  a  silver-white,  lustrous  metal,  of  specific  gravity 
8.9.  It  melts  at  about  1485°.  Nickel  is  malleable,  ductile, 
and  tenacious.  It  is  but  slightly  altered  on  exposure  to  the 
air,  and  consequently  it  is  frequently  used  in  plating  other 
metals.  Nickel  is  but  slowly  attacked  by  sulphuric  or  hydro- 
chloric acid ;  but  nitric  acid  dissolves  it  readily.  Like  iron, 
nickel  is  magnetic,  and  exhibits  the  phenomena  of  passivity. 

Nickel  is  used  in  many  alloys.  German  silver,  or  argentan,  an 
alloy  of  brass  with  nickel,  has  already  been  mentioned.  Man- 
ganine,  an  alloy  of  nickel  with  copper  and  manganese,  is  used 
for  resistance  wires.  Nickel  coins  contain  75  per  cent  copper 
and  only  25  per  cent  nickel,  which  fact  demonstrates  the  great 
power  of  nickel  to  impart  its  color  to  alloys.  About  three  per 
cent  nickel  added  to  steel  produces  a  product  of  great  strength; 
and  much  nickel  is  used  in  making  nickel  steel  for  armor 
plates. 

The  Chinese  have  used  nickel-copper  alloys  under  the  name 
packfong  for  many  centuries.  Through  the  work  of  Cronstedt 
and  Bergmann,  nickel  was  distinguished  from  other  metals  in 
1751.  Nickel  coins  came  into  use  about  fifty  years  ago. 

Nickel  Oxides  and  Hydroxides.— Nickelous  hydroxide  Ni(OH)2 
is  an  apple-green,  amorphous  precipitate,  formed  by  adding 


IRON,   NICKEL,   AND   COBALT  517 

caustic  alkali  to  a  solution  of  a  nickel  salt.  On  ignition  it 
yields  a  green  powder,  nickelous  oxide  NiO.  Nickelic  oxide 
Ni2O3  is  a  black  powder  obtained  by  careful  ignition  of  the 
nitrate.  Nickelic  hydroxide  Ni(OH)3  is  formed  when  a  solu- 
tion of  a  nickel  salt  is  treated  with  an  alkaline  hypochlorite. 
It  is  a  black  precipitate. 

Salts  of  Nickel.  — In  these  nickel  is  always  bivalent.  The  an- 
hydrous salts  are  yellow  or  brown,  and  the  hydrous  salts  and 
the  solutions  are  green.  Nickelic  oxide  acts  like  a  peroxide 
on  treatment  with  acids ;  there  are  no  nickelic  salts. 

Nickelous  chloride  NiCl2-6H2O  forms  green,  monoclinic 
srystals  that  readily  dissolve  in  water. 

Nickelous  nitrate  Ni(NO3)2-6  H2O  consists  of  green,  deli- 
quescent monoclinic  plates. 

Nickelous  sulphate  NiSO4  •  7  H2O  forms  green,  orthorhombic 
prisms  that  are  isomorphous  with  the  other  vitriols.  It  may 
also  be  obtained  as  NiSO4  •  6  H2O  in  form  of  tetragonal  crystals. 
It  is  readily  soluble  in  water.  The  crystals  effloresce  on  ex- 
posure to  the  air.  With  ammonium  sulphate,  nickel  sulphate 
forms  nickel  ammonium  sulphate  (NH4)2SO4-NiSO4-6  H2O, 
which  is  isomorphous  with  Mohr's  salt.  Nickel  ammonium 
sulphate  is  soluble  in  about  17  parts  of  water.  It  is  used  in 
nickel  plating,  in  which  process  a  thick  nickel  plate  is  used  as 
anode,  and  the  thoroughly  cleaned  object  to  be  plated  is  the 
cathode. 

Nickelous  cyanide  Ni(CN)2  is  formed  as  a  green  precipitate 
by  adding  potassium  cyanide  to  a  solution  of  a  nickel  salt.  In 
excess  of  potassium  cyanide,  the  precipitate  dissolves,  forming 
Ni(CN)2  -  2  KCN  •  H2O,  which  may  be  obtained  as  reddish  yellow 
monoclinic  crystals.  On  boiling,  the  solution  remains  un- 
changed ;  but  on  treatment  with  hypochlorites,  black  nickelic 
hydroxide  is  precipitated,  which  fact  is  used  in  separating 
nickel  from  cobalt. 

Nickelous  sulphide  NiS  is  formed  as  a  black  precipitate  when 
alkali  sulphides  are  added  to  solutions  of  nickel  salts.  The 
sulphide  is  slightly  soluble  in  excess  of  alkali  sulphides,  yielding 
a  dark  brown  solution  from  which,  by  addition  of  acetic  acid, 
the  sulphide  is  again  precipitated.  Nickelous  sulphide  dissolves 
but  slightly  in  dilute  hydrochloric  acid. 

Nickel  carbonyl  Ni(CO)4  is  a  colorless  liquid  boiling  at  43° 


518  OUTLINES   OF  CHEMISTRY 

and  congealing  at  —  25°.  It  is  formed  by  passing  carbon  mo- 
noxide over  nickel  obtained  by  reducing  the  oxide  or  oxalate 
by  ignition  in  hydrogen.  The  vapors  of  nickel  carbonyl  are 
poisonous.  The  substance  burns,  depositing  nickelous  oxide 
and  nickel.  This  compound  is  similar  to  iron  carbonyl  Fe(CO)5, 
which  is  formed  by  passing  carbon  monoxide  over  finely  divided 
iron  under  a  pressure  of  about  six  atmospheres  at  40°  to  80°. 
An  analogous  compound  of  cobalt  has  not  been  made. 

Occurrence,  Preparation,  and  Properties  of  Cobalt. — Cobalt  is 
found  in  meteorites  and  in  smaltite  CoAs2  and  cobaltite  CoAsS, 
in  which  it  is  commonly  associated  with  nickel,  iron,  and  man- 
ganese. The  principal  localities  are  Canada,  Sweden,  Bohemia, 
Germany,  and  the  Urals. 

The  metal  is  obtained  by  igniting  the  oxides  or  oxalate  in  a 
current  of  hydrogen.  Cobaltous  oxide  is  also  reduced  by  mix- 
ing it  with  starch  or  flour  and  making  small  cubes  of  the  paste 
formed,  which  are  then  embedded  in  pulverized  carbon  and 
highly  ignited  in  a  crucible.  Thus  the  oxide  is  reduced,  and 
the  metal  is  obtained  in  form  of  cubes.  Metallic  nickel  may 
be  obtained  similarly.  Cobaltous  oxide  may  be  reduced  by  the 
Goldschmidt  process. 

Cobalt  is  a  silver- white,  malleable,  tenacious  metal  of  specific 
gravity  8.5.  It  melts  at  about  1500°.  On  exposure  to  the 
air,  it  soon  acquires  a  reddish  hue.  Red-hot  cobalt  decomposes 
steam.  The  metal  is  magnetic.  It  dissolves  in  nitric  acid, 
but  other  dilute  acids  attack  it  but  slowly.  The  atomic  weight 
of  cobalt  is  58.97  and  its  valence  is  either  two  or  three.  The 
cobaltous  compounds  are  the  more  common.  The  metal  itself 
is  not  used  in  the  arts,  but  its  oxides  are  used  to  color  glass  and 
porcelain  blue.  The  fact  that  silicates  are  thus  colored  by 
cobalt  ores  has  been  known  since  ancient  times,  but  the  metal 
was  not  obtained  till  1735,  when  the  Swedish  chemist  Brandt 
prepared  it.  The  word  cobalt  means  goblin.  This  name  was 
given  to  the  ores  of  cobalt  by  the  miners ;  for  though  these 
ores  have  a  bright  metallic  luster  suggesting  a  metallic  content, 
it  was  not  till  relatively  recent  times  that  the  metal,  was  ob- 
tained from  them. 

Oxides  and  Hydroxides  of  Cobalt.  —  Cobaltous  hydroxide 
Co(OH)2  is  a  rose-red  precipitate  formed  by  adding  caustic 
alkali  to  a  solution  of  a  cobaltous  salt,  and  then  boiling  to 


IRON,   NICKEL,   AND   COBALT  519 

decompose  the  blue  basic  salt  that  is  first  precipitated.  On 
igniting  cobaltous  hydroxide  or  carbonate  out  of  contact  with 
the  air,  cobaltous  oxide  CoO  is  obtained  as  a  greenish  powder. 
This  is  also  formed  by  passing  steam  over  red-hot  cobalt.  On 
heating  cobaltous  oxide  in  the  air,  cobaltous  cobaltic  oxide  Co3O4, 
which  is  analogous  to  magnetite,  is  formed  as  a  black  powder. 
This  is  produced  as  the  final  product  by  igniting  any  oxide  of 
cobalt  in  the  air.  It  also  results  by  calcining  the  nitrate.  But 
if  the  latter  is  gently  heated,  cobalt  sesquioxide  Co2O3  is  first 
obtained  as  a  dark  brown  powder,  'which  yields  oxygen  and 
Co3O4  on  further  ignition.  Cobaltic  hydroxide  Co(OH)3  is  a 
dark  powder  formed  by  treating  a  solution  of  a  cobaltous  salt 
with  hypochlorites.  Cobalt  dioxide  CoO2  is  also  known. 

Other  Cobalt  Compounds.  —  The  hydrous  cobaltous  salts  are 
red,  which  is  also  true,  as  a  rule,  of  their  solutions  at  ordinary 
temperatures.  The  anhydrous  cobaltous  salts  'are  generally 
blue,  as  are  many  of  their  solutions  at  higher  temperatures. 
But  few  simple  cobaltic  salts  are  known. 

Cobaltous  chloride  CoCl2  is  a  blue  crystalline  powder  obtained 
by  heating  cobalt  in  chlorine.  •  From  aqueous  solutions,  it  may 
be  obtained  in  form  of  deep  red,  monoclinic  prisms  of  the  com- 
position CoCl2  •  6  H2O.  The  salt  is  also  formed  by  dissolving 
cobaltous  hydroxide  or  carbonate  in  hydrochloric  acid.  When 
the  hydrous  cobaltous  chloride  is  heated,  it  turns  blue,  becom- 
ing red  again  as  it  absorbs  water  after  cooling.  So  by  writing 
on  paper  with  a  solution  of  cobaltous  chloride  (so-called  sym- 
pathetic ink)  the  lines,  which  are  at  first  invisible,  become  blue 
on  gently  heating  the  paper.  On  exposure  to  the  air  at  ordi- 
nary temperatures,  the  salt  again  takes  on  water  and  the  writ- 
ing fades. 

Cobaltous  nitrate  Co(NO3)2  forms  red,  deliquescent,  mono- 
clinic  prisms. 

Cobaltous  sulphate  CoSO4-7H2O  is  isomorphous  with  the 
other  vitriols.  It  loses  water  readily  on  heating.  The  an- 
hydrous salt  is  red  and  is  not  easily  decomposed,  even  on 
strong  ignition. 

Cobaltic  sulphate  Co2(SO4)3-18  H2O  has  been  obtained  at 
the  anode  by  electrolysis  of  cobaltous  sulphate. 

Cobalt  silicates  result  when  glass  or  silica  and  potash  are 
fused  with  an  oxide  or  other  compound  of  cobalt.  The  blue 


520  OUTLINES   OF   CHEMISTRY 

glass  that  results  is  pulverized  and  used  as  a  pigment  under 
the  name  smalt.  It  serves  well  in  decorating  porcelain  and 
glass,  where  it  is  afterwards  fired  in ;  but  it  is  not  satisfactory 
in  oil  or  water  paints. 

Cobaltous  sulphide  CoS  is  obtained  as  a  black  precipitate  by 
adding  alkali  sulphides  to  solutions  of  cobalt  salts.  Like  the 
sulphide  of  nickel,  cobaltous  sulphide  is  insoluble  in  dilute 
hydrochloric  acid.  The  sulphides  Co2S3,  Co3S4,  and  CoS2  occur 
in  nature. 

Cobaltous  carbonate  CoCO3  is  a  bright  red  powder.  By  add- 
ing alkali  carbonates  to  solutions  of  cobaltous  salts,  red  pre- 
cipitates of  basic  carbonates  form,  which  decompose  and  turn 
blue  on  boiling. 

Cobalt  silicides  Co2Si  and  CoSi  have  been  made  in  the  electric 
furnace.  They  are  very  hard  crystalline  substances. 

Cobaltous  cyanide  Co(CN)2  is  formed  as  a  reddish  precipitate 
by  adding  potassium  cyanide  to  a  solution  of  a  cobaltous  salt. 
Cobaltous  cyanide  is  soluble  in  an  excess  of  potassium  cyanide, 
forming  potassium  cobaltous  cyanide  K4Co(CN)6,  which  upon 
oxidation  yields  potassium  cobaltic  cyanide  K3Co(CN)6.  These 
compounds  are  similar  to  those  of  the  corresponding  double 
cyanides  of  iron.  Nickel  does  not  form  analogous  double  cyanides. 

Potassium  cobaltic  nitrite  3  Co(NO2)3-6KNO2'3  H2O  is 
formed  as  a  yellow  precipitate  when  potassium  nitrite  is  added 
to  a  solution  of  a  cobaltous  salt  acidified  with  acetic  acid.  The 
precipitate  separates  out  slowly.  It  is  called  cobalt-yellow. 
The  reaction  is  used  to  separate  cobalt  from  nickel.  The  lat- 
ter does  not  form  an  analogous  salt. 

Cobalt  amines  are  complex  compounds  which  cobaltic  salts 
form  with  ammonia.  These  compounds  are  made  by  treating 
solutions  of  cobaltous  chloride  with  ammonia  in  sufficient 
amount  to  redissolve  the  cobaltous  hydroxide  that  first  forms. 
On  standing  in  the  air,  these  brown  solutions  absorb  oxygen, 
thus  becoming  red.  When  hydrochloric  acid  is  added  to  the 
red  solution,  brick-red  crystals  of  roseo  cobaltic  chloride 
CoCl3-5NH3-H2O  are  deposited.  On  boiling,  the  solution 
yields  purpureo- cobaltic  chloride  CoCl3(NH3)5,  which  forms  a 
bright  red  crystalline  powder,  or  brownish  luteo-cobaltic  chlo- 
ride CoCl3-6NH3,  if  much  ammonium  chloride  is  present  in 
the  strongly  acid  solution.  '  Other  cobaltous  salts  form  similar 


IRON,   NICKEL,    AND   COBALT  521 

complicated  compounds  on  treatment  with  ammonia.  The  three 
examples  given  must  suffice  here  to  indicate  their  general  char- 
acter, though  the  compounds  are  very  interesting  and  have  in 
recent  years  received  careful  study  in  connection  with  the  sub- 
ject of  valence,  particularly  by  A.  Werner  and  S.  M.  Jorgensen. 
Analytical  Tests  for  Nickel  and  Cobalt.  —  When  heated  in  the 
borax  bead,  cobalt  compounds  always  give  a  blue  coloration. 
Nickel  compounds  yield  a  brown  bead  in  the  oxidizing  flame, 
and  a  milky  bead  in  the  reducing  flame,  for  by  the  latter 
metallic  nickel  is  formed.  The  reactions  of  salts  of  nickel 
and  cobalt  with  alkali  hydroxides,  sulphides,  hypochlorites, 
cyanides,  and  nitrites,  as  above  explained,  are  used  in  the 
detection  and  separation  of  nickel  and  cobalt. 


CHAPTER  XXXI 

THE  METALS   OF   THE   PLATINUM   FAMILY 

THE  metals  of  the  platinum  group  and  their  symbols 
and  atomic  weights  are:  ruthenium  (Ru — 101.7),  rhodium 
(Rh— 102.9),  palladium  (Pd— 106.7),  osmium  (Os— 190.9), 
iridium  (Ir  — 193.1),  and  platinum  (Pt  — 195.2).  These 
metals  always  occur  together  in  nature  in  the  form  of  alloys  con- 
sisting of  small  corroded  metallic  bits  or  nuggets  found  in 
alluvial  sands,  chiefly  in  the  Urals,  California,  Brazil,  Australia, 
Borneo,  and  Sumatra.  These  alloys  contain  from  50  to  85  per 
cent  platinum,  several  per  cent  iridium,  and  generally  less  than 
two  per  cent  of  each  of  the  other  members  of  the  group.  Gold, 
copper,  iron,  and  other  metals  are  also  usually  present.  By 
careful  washing  of  the  sands,  the  particles  of  platinum  ore  are 
obtained  as  in  the  case  of  placer  mining  of  gold.  Though 
platinum  ore  usually  occurs  in  the  form  of  mere  grains,  larger 
nuggets  weighing  up  to  about  eight  kilograms  have  been  found 
in  rare  instances. 

It  will  be  observed  that  ruthenium,  rhodium,  and  palladium 
have  approximately  the  same  atomic  weight;  their  specific 
gravities  are  also  nearly  alike,  being  12.3,  12.1,  and  11.5,  re- 
spectively. Again,  the  second  three  metals,  osmium,  iridium, 
and  platinum,  have  nearly  the  same  atomic  weight,  which  is 
about  90  units  higher  than  that  of  the  first  three  members. 
Osmium,  iridium,  and  platinum  also  have  about  the  same  spe- 
cific gravities,  namely,  22.5,  22.4,  and  21.5,  respectively.  It  will 
be  recalled  that  the  atomic  weights  of  iron,  nickel,  and  cobalt 
are  also  nearly  the  same,  and  that  these  metals  have  specific 
gravities  that  differ  but  slightly  from  one  another. 

Extraction  of  Platinum  from  the  Ores.  The  ores  are  freed 
from  sand  and  other  adhering  impurities,  and  then  treated  with 
dilute  aqua  regia,  which  dissolves  the  gold,  copper,  and  iron 
that  may  be  present,  leaving  the  platinum  metals  behind.  On 
adding  concentrated  aqua  regia  to  this  residue,  platinum,  pal- 

522 


THE  METALS  OF  THE  PLATINUM  FAMILY       523 

ladium,  ruthenium,  rhodium,  and  some  of  the  iridium  are 
dissolved.  The  residue  left  consists  of  osmium  and  iridium 
together  with  some  ruthenium.  From  the  solution,  the  plati- 
num and  iridium  are  precipitated  by  means  of  ammonium  chlo- 
ride as  (NH4)2PtCl6  and  (NH4)2IrCl6.  There  is  usually  but 
little  iridium  present  in  this  precipitate.  On  ignition  of  the 
latter,  metallic  platinum,  alloyed  with  a  little  iridium,  is  ob- 
tained as  a  spongy  mass.  As  platinum  containing  minor 
amounts  of  iridium  is  excellent  for  crucibles,  dishes,  and  other 
purposes,  being  stronger  mechanically  and  more  resistant  to- 
wards reagents  than  pure  platinum,  this  sponge  is  commonly 
used  without  further  attempts  at  purification.  From  the  filtrate 
containing  palladium,  rhodium,  and  ruthenium,  these  metals  are 
precipitated  by  means  of  iron  and  then  separated  by  special 
processes  which  will  not  be  described  here. 

Ruthenium  (Ru — 101.7)  was  discovered  by  Glaus  in  1845. 
The  element  is  named  after  the  Ruthenians,  a  little  Russian 
people  in  whose  country  the  metal  is  found.  Ruthenium  occurs 
in  the  alloy  of  osmium  and  iridium,  osmiridium,  remaining  as 
a  residue  when  platinum  ores  are  treated  with  concentrated 
aqua  regia.  Ruthenium  is  commonly  prepared  from  this  alloy. 
It  is  a  brittle  metal  of  steel-gray  color.  It  melts  above  2000°. 
When  heated  in  the  air,  ruthenium  oxidizes.  The  oxides  RuO, 
Ru2O3,  RuO2,  and  RuO4  are  known.  They  are  black  powders. 
Ruthenium  is  not  attacked  by  acids ;  even  aqua  regia  acts  but 
slowly  on  it,  forming  the  trichloride  RuCl3.  The  dichloride 
RuCl2  is  known,  but  the  tetrachloride  RuCl4  exists  only  in 
solutions.  On  fusing  ruthenium  with  caustic  potash  and  salt- 
peter, potassium  ruthenate  K2RuO4,  readily  soluble  in  water,  is 
obtained,  which  fact  is  used  in  separating  ruthenium  from 
osmiridium. 

Rhodium  (Rh  — 102.9)  was  discovered  by  Wollaston  in  1803. 
It  receives  its  name  from  the  red  color  of  its  chloride.  The 
metal  is  silver- white,  ductile,  and  malleable.  It  melts  at  about 
2000°,  and  when  pure  it  is  not  attacked  by  acids,  not  even  by 
aqua  regia,  which,  however,  dissolves  rhodium  alloyed  with 
platinum.  By  fusing  rhodium  with  caustic  potash  and  saltpeter 
and  extracting  the  mass  with  nitric  acid,  oxides  of  rhodium  are 
obtained.  The  oxides  of  rhodium  are  RhO,  Rh2O3,  and  RhO2  ; 
the  hydroxides  Rh(OH)3  and  Rh(OH)4  are  also  known.  The 


524  OUTLINES   OF  CHEMISTRY 

red  chloride  RC13,  formed  by  the  action  of  chlorine  on  the  metal, 
is  the  only  one  isolated.  It  is  insoluble  in  water  and  acids.  It 
forms  soluble  double  salts  with  alkali  chlorides.  Rhodium 
chloride  forms  RhCl3  •  (NH3)5  and  other  compounds  analogous 
to  cobalt  amines.  The  double  cyanide  K3Rh(CN)6,  which  is 
analogous  to  K3Fe(CN)6,  is  also  known. 

Palladium  (Pd  — 106.7)  was  also  discovered  by  Wollaston 
in  1803.  It  is  named  from  Pallas,  an  asteroid  discovered  in 
1802.  While  always  present  in  platinum  ores,  palladium  also 
occurs  in  alloys  with  gold  in  Brazil.  Silver  generally  contains 
traces  of  palladium.  It  is  a  silver-white,  ductile,  malleable 
metal,  which  melts  at  about  1535°.  Palladium  absorbs  over  300 
times  its  volume  of  hydrogen  under  ordinary  conditions,  at  100° 
about  twice  as  much  hydrogen  is  absorbed.  On  electrolyzing 
water,  using  a  palladium  cathode,  the  latter  absorbs  nearly 
1000  times  its  volume  of  hydrogen.  The  compound  formed, 
palladium-hydrogen,  is  lighter  than  the  metal ;  its  composition 
varies  with  the  conditions  under  which  it  is  prepared.  It  gives 
off  its  hydrogen  completely  when  heated  to  redness,  though 
evolution  of  the  gas  begins  even  at  100°.  The  hydrogen 
liberated  effects  reductions  like  nascent  hydrogen. 

Palladium  is  soluble  in  nitric  acid,  in  aqua  regia,  and  in  hot 
sulphuric  acid.  The  oxides  are  PdO  and  PdO2 ;  both  are 
black.  The  dioxide  yields  the  chloride  PdCl2  and  chlorine, 
when  treated  with  hydrochloric  acid,  PdCl4  being  unknown. 
The  following  typical  salts  are  also  well  known :  — 

Pd(N03)2;   PdI2;   PdS04;   K2PdCl6;   PdCl2  -  (NH8)a. 

Osmium  (Os  — 190.9)  was  discovered  in  1804  by  Tennant. 
It  is  found  in  platinum  ores,  in  osmiridium  and  iridosmium, 
two  alloys  of  osmium  and  iridium. 

Osmium  is  a  steel-gray,  hard,  brittle  metal,  which  has  the 
highest  specific  gravity  of  all  known  substances,  22.5.  It  melts 
in  the  electric  furnace  above  2500°.  Osmium  is  oxidized  to 
OsO4  by  aqua  regia,  nitric  acid,  or  by  ignition  in  the  air. 
Osmium  tetroxide  forms  white,  monoclinic  crystals  that  read- 
ily melt,  and  then  boil  at  about  100°,  emitting  a  very  pungent 
vapor,  from  whose  unbearable  odor  osmium  gets  its  name. 
Osmium  tetroxide  is  used  in  hardening  and  staining  histological 
specimens,  being  commonly  called  osmic  acid  or  perosmic  acid. 


THE   METALS    OF  THE   PLATINUM  FAMILY  525 

Its  use  depends  upon  the  fact  that  in  contact  with  organic 
tissues,  especially  with  fats,  it  is  reduced,  forming  finely 
divided  metallic  osmium.  The  readiness  with  which  osmium 
tetroxide  is  formed  and  its  volatility  serve  in  separating  os- 
mium from  iridium  and  ruthenium,  which  also  occur  in  osmi- 
ridium.  The  chlorides  OsCl2,  Os2Cl6,  and  OsCl4  are  known,  as 
are  also  osmates  like  K2OsO4-2H2O.  The  oxides  are  OsO, 
Os2O8,  OsO2  and  OsO4. 

Iridium  (Ir  — 193.1)  was  also  discovered  by  Tennant  in  1804. 
It  is  a  grayish  white,  rather  brittle  metal  which  melts  at  about 
2200°.  It  is  found  in  ores  of  platinum  and  in  osmiridium.  Its 
separation  from  osmium  has  been  mentioned  in  connection  with 
the  latter.  From  platinum  it  may  be  separated  by  using  the 
fact  that  the  compound  (NH4)2IrCl6  is  readily  soluble  in 
water.  Aqua  regia  attacks  pure  iridium  but  slowly.  The 
oxides  are  IrO,  Ir2O3,  and  IrO2.  The  hydroxides  Ir(OH)8  and 
Ir(OH)4  and  the  chlorides  IrCl2,  IrCl8,  and  IrCl4  are  also 
known.  Iridium  gets  its  name  from  the  fact  that  it  forms 
compounds  of  various  colors. 

Platinum  (Pt  — 195.2)  is  a  silver- white,  tenacious,  malleable, 
and  ductile  metal  of  specific  gravity  21.47.  The  name  plati- 
num comes  from  platina,  which  is  a  diminutive  of  the  Spanish 
plata,  meaning  silver.  Platinum  melts  in  the  oxyhydrogen 
flame  at  about  1777°.  At  white  heat  it  may  be  welded  with- 
out the  aid  of  a  flux,  for  its  surface  is  not  covered  with  oxide. 
Even  at  high  temperatures,  platinum  does  not  oxidize  in  the 
air.  Aqua  regia  dissolves  platinum,  but  the  latter  is  not 
attacked  by  hydrofluoric,  hydrochloric,  sulphuric,  or  nitric 
acid.  Molten  caustic  alkalies  and  alkali  nitrates,  cyanides,  or 
sulphides  attack  platinum,  and  it  readily  forms  alloys  with  most 
of  the  heavy  metals  and  also  with  silicon,  boron,  phosphorus,  ar- 
senic, and  antimony.  These  facts  are  of  importance,  for  the 
substances  mentioned,  and  also  such  compounds  as  may  yield  these 
substances  on  ignition,  must  not  be  heated  in  platinum  dishes. 
Platinum  should  also  not  be  heated  in  a  reducing  flame,  for  thus 
the  metal  takes  up  carbon  and  becomes  brittle. 

Platinum  sponge  is  formed  by  ignition  of  ammonium  platinic 
chloride;  while  platinum  black  is  produced  by  electrolysis  of 
platinic  chloride  solutions,  or  by  adding  finely  divided  magne- 
sium, iron,  or  zinc  to  the  latter.  In  finely  divided  form,  platinum 


526 


OUTLINES   OF   CHEMISTRY 


absorbs  about  300  times  its  volume  of  hydrogen  or  about  one  third 
that  volume  of  oxygen.  The  gases  are  said  to  be  occluded  in 
the  metal.  On  ignition  they  escape.  The  catalytic  oxidations 
and  reductions  effected  by  finely  divided  platinum  depend 

upon  the  fact  that  the  latter  absorbs 
the  gases  mentioned,  which  then  act 
far  more  vigorously  than  when  in  the 
ordinary  state.  Thus  Dobereiner's 
lamp  (Fig.  155)  depends  upon  the 
fact  that  when  a  jet  of  hydrogen  is 
directed  against  a  platinum  sponge 
the  latter  is  heated  to  redness  and  so 
lights  the  jet.  The  decomposition  of 
hydrogen  peroxide  and  the  synthesis 
of  SO3  from  SO2  and  O2  are  further 
typical  instances  of  the  catalytic  action 
FlG  155  of  platinum  black. 

Besides   being   used    for    crucibles, 

dishes,  and  other  utensils  for  chemical  operations,  much  plati- 
num is  required  for  making  various  electrical  connections. 
Platinum  salts  are  used  in  photography. 

There  are  two  series  of  platinum  compounds.  In  platinous 
compounds  the  metal  is  bivalent,  while  in  platinic  compounds  it 
is  quadrivalent.  Platinic  chloride  PtCl4  is  formed  by  the 
action  of  chlorine  on  platinum  at  high  temperatures,  or  by 
ignition  of  chlorplatinic  acid  H2PtCl6  •  6  H2O,  which  results 
when  platinum  is  dissolved  in  aqua  regia.  With  the  exception 
of  sodium  chlorplatinate  Na2PtCl6  the  alkali  salts  of  chlor- 
platinic acid  are  insoluble  in  alcohol  and  sparingly  soluble  in 
water,  hence  the  use  of  chlorplatinic  acid  in  separating  sodium 
from  potassium  in  analysis.  Potassium  chlorplatinate  K2PtCl6,  or 
PtCl4.2KCl,  forms  a  golden  yellow,  crystalline  precipitate. 
The  crystals  are  small  octahedra;  they  are  isomorphous  with 
ammonium  chlorplatinate  (NH4)2PtCl6.  When  potassium  chlor- 
platinate is  ignited,  potassium  chloride  and  platinum  remain ; 
whereas  when  ammonium  chlorplatinate  is  heated,  the  residue 
consists  simply  of  spongy  platinum.  Either  of  these  salts, 
which  are  also  called  potassium  platinic  chloride  and  ammo- 
nium platinic  chloride,  is  readily  formed  by  adding  the  alkali 
chloride  to  a  solution  of  chlorplatinic  acid  or  platinic  chloride. 


THE   METALS  OF  THE  PLATINUM  FAMILY  527 

On  heating  chlorplatinic  acid  to  300°,  or  on  passing  chlorine 
over  platinum  sponge  at  about  245°,  platinous  chloride  PtCl2  is 
obtained  as  a  grayish  green,  insoluble  powder,  which  on  igni- 
tion leaves  platinum  as  a  residue.  Potassium  platinous  chloride 
K2PtCl4  is  formed  by  reducing  potassium  platinic  chloride 
with  cuprous  chloride.  Platinbus  hydroxide  Pt(OH)2  is  formed 
by  treating  platinous  chloride  with  caustic  alkali.  On  ignition 
of  platinous  hydroxide,  platinous  oxide  PtO  is  formed,  and 
finally  platinum.  Platinic  hydroxide  Pt(OH)4  results  when 
platinic  chloride  is  treated  with  caustic  alkali.  By  careful 
ignition  of  this  hydroxide,  platinic  oxide  may  be  obtained, 
which  on  strong  ignition  yields  platinum.  When  platinic 
hydroxide  is  treated  with  an  excess  of  caustic  alkali,  platinates 
are  formed,  thus  :  — 

Pt(OH)4  +  4  NaOH  =  Na4PtO4  +  4  H2O. 

Platinates  also  result  when  platinum  is  placed  in  molten  caustic 
alkalies. 

Platinous  sulphide  PtS  and  platinic  sulphide  PtS2  are  black 
precipitates,  insoluble  in  acids,  formed  by  adding  hydrogen 
sulphide  to  solutions  of  platinous  and  platinic  compounds, 
respectively.  These  sulphides  are  soluble  in  aqua  regia. 
They  also  dissolve  in  alkali  sulphides,  with  which  they  form 
sulpho-salts. 

Analytical  Tests  for  Platinum.  —  The  sulphides  described  are 
characteristic.  On  heating  any  platinum  compound  with  soda 
on  charcoal,  spongy  platinum  is  obtained.  This  is  soluble  in 
aqua  regia,  and  from  the  solution  potassium  chlorplatinate  may 
be  precipitated  by  adding  potassium  chloride. 


INDEX 


Abietic  acid,  264 

Absorption  spectrum,  368,  369 

of  blood,  370 
Acetates,  244 
Acetic  acid,  242,  243 

glacial,  244 
Acetone,  255 
Acetylene,  235 
Acheson  graphite,  212 
Acid,  definition  of,  121,  125,  430 
Acidimetry,  128 
Acids,  120 

anhydrides  of,  122 

basicity  of,  126 

strength  of,  137 
Acker  process,  357 
Actinium,  392 
Adsorption,  19,  214 
Affinity,  chemical,  11,  134,  135 
Agate,  291 
Air,  143 

alkaline,  150 

ammonia  in,  145 

bacteria  and  microbes  in,  144,  146 

carbon  dioxide  in,  145 

composition  of,  143,  144 

fixed,  139,  223 

humidity  of,  36 

liquid,  147 

mephitic,  139 

moisture  in,  144 

nature  of,  143 

phlogisticated,  139 
Alabaster,  380 
Albite,  294,  465 
Albuminoids,  264 
Albumins,  264 
Alchemists,  194 
Alcohols,  237,  238 
Aldebaranium,  470 
Aldehydes,  240,  241 
Algarotus,  326 
Algin,  113 
Alkalies,  343 
Alkali  metals,  343,  365 
Alkalimetry,  128 
Alkaline  earth  metals,  374 

detection  of,  390 
Alkaloids,  263 


Allotropism  (see  Allotropy) 
Allotropy,  89,  208 
Alloys  of  copper,  439 

of  manganese,  496 
Alumina,  460 
Aluminates,  461 
Aluminum,  457 

acetate,  464 

alloys  of,  459 

bromide,  463 

bronze,  440 

chloride,  463 

fluoride,  463 

hydroxide,  460 

iodide,  463 

oxide,  460 

production  of,  458 

properties  of,  458 

silicates,  465 

sulphate,  463 

sulphide,  463 

tests  for,  467 

uses  of,  459 
Alums,  464 
Alunite,  464 

Amalgamation  process,  445,  452 
Amalgams,  402 

for  filling  teeth,  474 
Amethyst,  291 
Amines,  262 
Ammonia,  150 

action  of,  on  metals,  155,  156 

combustion  of,  in  oxygen,  154 

composition  of,  152,  153 

concentrated,  154 

liquid,  156 

of  crystallization,  156 

preparation  of,  151 

properties  of,  152 

soda  process,  360 

water,  154 
Ammonium,  154 

alum,  464 

amalgam,  403 

bichromate,  489 

bisulphate,  371 

bromide,  371 

carbamate,  372 

carbonate,  372 

carbonate,  acid,  372 


529 


530 


INDEX 


Ammonium  —  Continued 

chloride,  371 

chlorplatinate,  526 

hydrazoate,  159 

hydrosulphide,  371 

iodide,  371 

molybdate,  491 

nitrate,  372 

nitrite,  372 

persulphate,  371 

phosphomolybdate,  491 

platinic  chloride,  373 

salts,  155,  370 

salts,  detection  of,  157,  372 

sulphantimonate,  329 

sulphantimonite,  329 

sulpharsenate,  323 

sulpharsenite,  323 

sulphate,  371 

sulphide,  371 

tartrate,  acid,  373 
Amorphous  substances,  45 
Amyl  acetate,  252 
Anaxagoras,  68 
Andalusite,  465 
Anglesite,  478 
Anhydrite,  380 
Aniline,  262 

dyes,  262 

hydrochloride.  262 
Anions,  425 
Anode,  425 
Anorthite,  465 
Anthracite,  215 
Antichlorine,  363 
Antimonic  acid,  328 
Antimonious  acid,  327 
Antimonium,  Triumphal  chariot  of,  328 
Antimony,  323 

cinnabar,  330 

compounds  of,  with  halogens,  325 

compounds  of,  with  sulphur,  329 

nitrate,  328 

oxides  and  oxy-acids,  327 

oxychlorides,  326 

pentachloride,  326 

pentafluoride,  327 

pentasulphide,  329 

pentiodide,  327    ' 

pentoxide,  329 

sulphate,  328 

tetroxide,  328 

tribromide,  327 

trichloride,  325 

trifluoride,  326 

triiodide,  327 

trioxide,  327 

trisulphide,  329 
Antimonyl  group,  327 

nitrate,  328 

sulphate,  328 
Apatite,  304 


Aqua  fortis,  161 

regia,  165 
Arabite,  238 
Aragonite,  376 
Argentan,  516 
Argentic  nitrate,  449 
Argentite,  444 
Argentum,  447 
Argol,  250 
Argon,  147,  148 
Argyrodite,  472 
Aromatic  series,  236 
Arrhenius,  428,  429 
Arrhenius's  theory,  428,  431,  432 
Arsenic,  317 

acid,  321 

compounds  of,  with  halogens,  320 

disulphide,  322 

iodides  of,  320 

oxides  and  oxy-acids  of,  320 

pentasulphide,  323 

pentoxide,  322 

tribromide,  320 

trichloride,  320 

trifluoride,  320 

trioxide,  320 

trisulphide,  322 

white,  320 
Arsenious  acid,  321 
Arsenites,  321 

Arseniureted  hydrogen,  318 
Arsenolite,  317 
Arsine,  318 
Asbestus,  394 
Assaying,  453 

Association  in  solution,  419 
Asymmetric  carbon  atom,  247 
Asymmetric  system,  185 
Atmosphere  (see  Air) 
Atomic  theory  of  matter,  64,  67 

volumes,  relation   to  atomic  weights, 
341 

weights,  determination  of,  74 

weights,  choosing  of,  79,  80 

weights,  table  of,  81 
Atoms,  65,  77 
Atropine,  263 
Aurates,  455 
Auric  chloride,  454 

compounds,  454 

oxide,  455 

sulphide,  455 
Aurous  chloride,  454 

compounds,  454 

cyanide,  455 

oxide,  455 

sulphide,  455 
Avogadro,  70,  74,  77 
Avogadro's  hypothesis,  70,  72,  153 
Azoimide,  159 
Azote,  139 
Azurite,  436,  444 


INDEX 


531 


Babbitt  metal,  479 
Baeyer,  95 
Balard,  105 
Band  spectrum,  368 
Barite,  387 
Barium,  387 

bromide,  388 

carbonate,  389 

chloride,  388 

chromate,  490 

compounds  of,  388 

dioxide,  388 

ferrate,  510 

fluoride,  38.8 

hydroxide,  388 

iodide,  388 

nitrate,  388 

oxide,  388 

sulphate,  389 

sulphide,  381,  389 

tellurate,  207 
Barley  sugar,  257 
Bases,  122,  125,  127,  137,  430 
Basic  open  hearth  process,  508 
Bauxite,  457,  460 
Becquerel,  390 
Becquerel  rays,  390 
Beer,  238 
Begasse,  258 
Bengal  lights,  387,  389 
Benzene,  236 
Benzine,  233 
Benzoic  acid,  244 
Bergmann,  516 
Berthelot,  279,  288 
Berthollet,  53 
Beryl,  393 
Beryllium,  393 
Berzelius,  40,  61,  62,  63,  70,  77,  80,  125, 

204,  223,  290,  299,  424,  470 
Bessemer  converter,  437,  507 

process,  507 
Bicarbonates,  219 
Bichromates,  488 

uses  of,  489 
Binary  compound,  76 
Bioses,  257 

Birkelund  and  Eyde  process,  163 
Bismuth,  330 

bromide,  331 

dichloride,  331 

dioxide,  332 

disulphide,  333 

fluoride,  331 

halogen  compounds  of,  331 

iodide,  331 

nitrate,  332 

oxides  of,  331 

oxy bromide,  331 

oxyfluoride,  331 


oxyiodide,  331 

oxynitrate,  332 

pentoxide,  332 

salts  of  oxy-acids,  332 

subnitrate,  332 

tetroxide,  332 

trioxide,  331 

trisulphide,  333 
Bismuthyl  sulphate,  332 
Bitter  almonds,  oil  of,  241 
Black,  Joseph,  223 
Black  ash,  359 
Black-jack,  396,  398 
Blast  flame,  272 
Blast  furnace,  503 
Blast  furnace  slag,  503 
Bleach,  382 

Bleaching  powder,  98,  99,  382 
Blood  charcoal,  214 
Blowpipe  flame,  272 
Blue  cup  battery,  434 
Blue  printing,  515 
Boiling  points  of  solutions,  41 T 
Bonds,  83,  84,  235,  236 
Bone  black,  214 
Borax,  302 
Borax  glass,  302 
Bordeaux  mixture,  443 
Boric  acid,  300 
Boric  anhydride,  302 
Boron,  300 

carbide,  303 

chloride,  303 

fluoride,  303 

hydride,  302 

nitride,  303 

sulphide,  303 
Bort,  210 
Bottger,  308 
Bouronite,  478 
Boussingault,  144 
Boyle,  304 
Brandt,  304,  518 
Brandy,  238 
Brass,  398,  439 
Bricks,  466 
Brin's  process,  25 
Britannia  metal,  324,  474 
British  thermal  unit,  276 
Bromates,  110 
Bromic  acid,  109,  110 
Bromine,  105,  106 

oxy-acids  of,  109 

uses  of,  111 

water,  106 
Bromoform,  237 
Bronzes,  440,  474 
Brucine,  263 
Bruyn,  Lobry  de,  158 
Bullets,  479 

Bunsen,  353,  366,  369,  394 
Bunsen  burner,  371 


532 


INDEX 


Burnt  alum,  464 
Bussy,  394 
Butane,  232 
Butter,  254 

fat,  254,  412 

of  antimony,  326 

of  tin,  475 
Butyric  acid,  244 

G 

Cadaverine,  264 
Cadmium,  400 

bromide,  401 

chloride,  400 

compounds  of,  400 

hydroxide,  400 

iodide,  401 

nitrate,  401 

oxide,  400 

sulphate,  401 

sulphide,  401 
Caesium,  353 
Calamine,  396 
Calaverite,  452 
Calcite,  376 
Calcium,  375 

aluminate,  462 

bicarbonate,  219,  376 

bromide,  382 

carbide,  383 

carbonate,  376 

chloride,  381 

chromate,  490 

cyanamide,  383 

fluoride,  381 

hydride,  376 

hydroxide,  377 

iodide,  382 

metaphosphate,  305 

nitride,  376 

oxalate,  245 

oxide,  377 

phosphate,  382 

phosphide,  383 

silicate,  383 

silicide,  383 

sucrate,  258 

sulphate,  380 

sulphide,  381 

sulphite,  381 
Calomel,  404 
Calorie,  276 
Calorimeters,  276 
Cane  sugar,  257 
Cannizzaro,  77,  78,  79 
Caoutchouc  membranes,  423 
Caramel,  257 
Carats,  454 
Carbohydrates,  255 
Carbolic  acid,  240 
Carbon,  210 

allotropic  forms  of,  209 


amorphous,  213 

atom,  properties  of,  217 

atomic  weight  of,  75 

chemical  behavior  of,  215 

cycle  of,  222 

pyrophoric,  213 
Carbonado,  210 
Carbonates,  219 
Carbon  bisulphide,  227 
Carbon  bisulphide  furnace,  228 
Carbon  dioxide,  120,  218 

early  work  on,  223 

formula  of,  76,  216 

physiological  effects  of,  221 

properties  of,  220 

relations  of,  to  life,  222 

solid,  221 

uses  of,  221 

Carbonic  acid,  120,  218 
Carbon  monoxide,  223 

absorption  of,  226,  441 

formula  of,  76,  225 

physiological  effects  of,  227 

properties  of,  225 
Carbon  oxysulphide,  226 
Carbon  tetrachloride,  237 
Carbonyl  chloride,  226 

group,  255 

iron,  226 

nickel,  226 
Carborundum,  299 
Carboxyl  group,  242 
Carlisle,  424 
Carnallite,  344,  394 
Carnelian,  291 
Casciorolo,  389 
Cassiopeium,  470 
Cassiterite,  472,  476 
ast  iron,  gray,  white,  505 
atalytic  action  of  platinum,  526 
Catalytic  agents,  192 
Cathode,  425 
Cations,  425 
Caustic  soda,  357 
Cavendish,  13,  33,  35,  139,  161 
Celestite,  386 
Celluloid,  262 
Cellulose,  261 

nitrates  of,  261 

solvent  for,  439 
ement,  378 
Cementite,  515 
eric  hydroxide,  470 
erite,  469,  470 
Cerium,  299,  470 
Cerussite,  478 
Chalcedony,  291 
lhalcocite,  436 
!halcopyrite,  436 
halk,  376 

ihamber  crystals,  196 
Chameleon  mineral,  499 


INDEX 


533 


Chameleon  solution,  499 
Chance,  359 
Chancel,  308 
Chaptal,  139 
Charcoal,  214 
Chemical  change,  1,  2,  5 

cause  of,  11 

factors  affecting  same,  12,  135 

rate  of,  12,  23 

types  of,  9 
Chemical  compound,  distinguished  from 

solution,  6 
Chemical  elements,  6 

distribution  of,  8,  9 

Chemical  formula,  interpretation  of,  82 
Chemical  reactions  in  electrolytes,  431 

in  insulators,  431 
Chemistry,  branches  of,  4,  5 

organic,  217 

scope  of,  1 
Chevreul,  253 

Chili  saltpeter,  113,  164,  361 
Chloral,  241 
Chloral  hydrate,  241 
Chlorates,  99 
Chlorauric  acid,  454 
Chloric  acid,  99 
Chlorination  process,  492 
Chlorine,  52 

action  of,  on  water,  55,  136 

bleaching  with,  55,  99 

compounds  with  oxygen,  56,  96 

dioxide,  56,  96 

elementary  nature  of,  121 

heptoxide,  97,  101 

monoxide,  96 

peroxide,  56,  96 

preparation  of,  53 

properties  of,  54 

reactions  of,  87 

uses  of,  55 
Chloroform,  237 
Chlorous  acid,  101 
Chlorplatiriic  acid,  526 
Chlorsulphonic  acid,  203 
Choke  damp,  274 
Chromates,  488 
Chrome  alums,  488 

green,  486 

iron  ore,  485 

steel,  486 

yellow,  490 
Chromic  acid,  490 

chloride,  487 

hydroxide,  486 

oxides,  486 

sulphate,  488 
Chromite,  485 
Chromites,  487 
Chromium,  485 

analytical  tests  for,  491 

properties  of,  486 


Chromium  trioxide,  488,  490 
Chromous  chloride,  487 

hydroxide,  486 

sulphate,  487 

Chromyl  chloride,  486,  490 
Chrysoberyl,  393,  462 
Cinchonine,  263 
Cinnabar,  401,  407 
Citric  acid,  251 
Clarke,  F.  W.,  8 
Claus,  523 

Clausius,  427,  428,  429 
Clausius's  electrolytic  theory,  427 
Clays,  465 
Cleve,  469 
Coal,  bituminous,  hard,  soft,  215 

gas,  265,  266 

tar,  uses  of,  265 
Cobalt,  502,  518 

amines,  520 

analytical  tests  for,  521 

compounds,  519 

dioxide,  519 

sesquioxide,  519 

silicides  of,  520 

yellow,  520 
Cobaltic  hydroxide,  519 

silicate,  519 

sulphate,  519 
Cobaltite,  317,  518 
Cobaltous  carbonate,  520 

chloride,  519 

cobaltic  oxide,  519 

cyanide,  520 

hydroxide,  518 

nitrate,  519 

oxide,  519 

sulphate,  519 

sulphide,  520 
Cocaine,  263 
Codeine,  263 
Colemanite,  302 
Collodion,  262 
Colloidal  solution,  295 
Colloids,  295 
Columbites,  336 
Columbium,  336 
Combining  weights,  61 

table  of,  63 

unit  of,  62 
Combustion,  26,  269 

earlier  views  of,  32 

heat  of,  28,  282 

in  air,  27 

of  oxygen  in  hydrogen,  32 

temperature  of,  27 
Compounds,  denned,  9,  10,  132 

stability  of,  288 
Concrete,  380 
Conductors  of  the  first  class,  423 

of  the  second  class,  423 
Condy's  disinfecting  fluid,  500 


534 


INDEX 


Confieldite,  472 
Congo  red,  130 
Consolute  liquids,  411 
Consolute  pairs,  412 
Copper,  436,  438 

analytical  tests  for,  444 

arsenite,  444 

carbonates,  443 

extraction  of,  from  ores,  437 

ferro cyanide,  442 

ferrocyanide  membrane,  421 

glance,  436 

in  plants  and  animals,  436 

matte,  437 

nitrate,  443 

pyrites,  436 

refining,  438 

salts  of  oxy-acids,  442 

sulphate,  442 

sulphate,  basic,  443 

sulphate,  double  salts  with,  443 
Copperas,  512 
Corpuscles,  432 
Corrosive  sublimate,  405 
Corundum,  460 
Cotton-seed  oil,  252 
Courtois,  111 
Crawford,  390  . 
Cream  of  tartar,  250 
Creosote,  240 
Crocoisite,  478,  485 
Cronstedt,  516 
Crookes,  468 
Crookesite,  468 
Cryolite,  102,  457 
Crystalline  substances,  45 
Crystalloids,  295 
Crystals,  180 
Crystal  systems,  180 
Cullinan  diamond,  211 
Cupellation,  446 
Cup  grease,  234 
Cuprammonium  sulphate,  443 
Cupric  ammonium  chloride,  441 

bromide,  442 

chloride,  441 

chloride,  hydrolysis  of,  132,  133 

compounds,  436 

cyanide,  442 

fluoride,  442 

hydroxide,  441 

oxide,  440 

sulphide,  444 
Cuprous  bromide,  441 

chloride,  441 

compounds,  436 

cyanide,  442 

fluoride,  442 

iodide,  441 

oxide,  440 

sulphide,  444 
Curie,  M.  and  Mme.,  391,  392 


Curtius,  151,  159 
Cyanates,  230 
Cyanic  acid,  231 
Cyanide  process,  452 
Cyanides,  229 

of  copper,  442 

of  iron,  513 
Cyanogen,  229 
Cymogene,  233 


Dalton,  59,  61,  66,  67,  68,  69,  71,  74 

Davy,  53,  111,  274,  344,  354,  424 

Deacon  process,  53 

Debierne,  392 

Decomposition,  double,  9,  415,  430 

hydrolytic,  131 

Definite  proportions,  law  of,  4,  58 
Delafontaine,  471 
d'Elhujar  brothers,  493 
Deliquescence,  46 
Del  Rio,  335 
Democritus,  68 
Desiccator,  46 
Developers,  449 

Deville,  H.  Sainte-Claire,  52,  137,  166 
Dextrine,  259,  261 
Dextrose,  256 
Dewar,  103,  146 
Dialysis,  294,  295 
Diamond,  210 
Diaspore,  460 
Diastase,  258,  259 
Didymium,  471 
Dilution,  heat  of,  281 
Dimorphism,  179,  185 
Dioxogen,  95 

Diphosphorus  tetraiodide,  312 
Dissociation,  109,  137 

in  solutions,  419 
Disthen,  465 

Distillation,  destructive,  150,  215 
Disulphuric  acid,  201 
Dobereiner,  337 
Dobereiner' s  lamp,  526 
Dolomite,  376,  394 
Dulong,  40,  79,  160 
Dulong  and  Petit,  law  of,  78 
Dumas,  40,  144,  223 
Dutch  metal,  439' 
Dutch  process,  483 
Dyeing  of  cotton,  462 

of  wool  and  silk,  463 
Dynamite,  254 
Dysprosium,  471 

E 

Earth,  infusorial,  293 
Earthenware,  466 
Earths,  alkaline,  374 

metals  of,  455 
Efflorescence,  46 


INDEX 


535 


Eka-aluminum,  340,  467 

Eka-boron,  340,  369 

Eka-silicon,  340,  472 

Elastine,  264 

Electric  batteries,  432,  434 

Electric  furnaces,  212 

Electrochemical  series  of  the  metals,  435 

Electrodes,  425 

Electrolysis,  423,  425 

Electrolytes,  423,  424 

migration  in,  429 
Electrolytic  dissociation,  428 

process,  483 

soda  process,  361 

solution  tensions,  435 

theories,  426 
Electromotive  force,  433 
Electrons,  432 
Electron  theory,  432 
Elements,  acid-forming,  123,  125 

base-forming,  123,  125 

chemical,  list  of,  7 

classification  of,  337 

groups  of,  7,  339 

transmutation  of,  392 
Emanations  from  radium,  etc.,  392 
Emerald,  393 
Emery,  460 
Emulsin,  259 
Emulsion,  412 
Enantiomorphism,  251 
Energy,  1 

conservation  of,  11 

electrical,  433 

free,  281 

total,  281 

transformation  of,  11 
Ending  ide,  76 
Endings  and  prefixes,  102 
Endothermic  changes,  275 
Enzymes,  259 
Epicurus,  68 
Epsom  salt,  46,  396 
Equations,  chemical,  76,  85 
Equilibrium,  chemical,  132,  134,  153 
Equivalents,  chemical,  22,  61,  67 
Erbium,  471 
Erythrite,  238 
Esters,  252 
Ethane,  232,  235 
Ethers,  254 
Ethyl  amine,  262 

borate,  301 

chloride,  237 

ether,  254 

nitrate,  252 

nitrite,  252 

silicate,  299 

sulphide,  255 
Ethylene,  235 

bromide,  235 
Europium,  471 


Euxenite,  469 
Exothermic  changes,  275 


Facts  relative  to  laws  and  theories,  68 
Faraday,  424,  425,  426,  429 
Faraday's  law,  426 

view  on  electrolysis,  427 
Fats,  252 

solubility  of,  47,  412 
Fatty  series,  236 
Fehling's  solution,  256 
Fermentation  and  ferments,  259 
Ferric  acetates,  515 

acid,  510 

ammonium  alum,  465 

bromide,  511 

chloride,  511 

ferrocyanide,  514 

hydroxide,  509 

nitrate,  515 

oxide,  509 

oxides,  hydrated,  510 

phosphate,  515 

sulphate,  513 

sulphide,  512 
Ferricyanic  acid,  514 
Ferrochromium,  486 
Ferrocyanic  acid,  514 
Ferromanganese,  496 
Ferrous  ammonium  sulphate,  513 

bicarbonate,  513 

bromide,  511 

carbonate,  513 

chloride,  510 

ferric  cyanide,  514 

ferric  oxide,  510 

hydroxide,  509 

iodide,  511 

nitrate,  515 

oxide,  509 

phosphate,  515 

sulphate,  512 

sulphide,  512 
Fertilizer  from  slag,  507 
Fertilizers,  198 
Fire  brick,  395 
Fire  damp,  274 
Fischer,  Emil,  264 
Flame,  268 

hydrogen,  20 

luminosity  of,  270 

oxidizing,  273  , 

reducing,  273 

reverse,  269 

singing,  20 

structure  of,  272 
Flash  light  powder,  394 
Flash  point,  233 
Flint,  291 
Flores  zinci,  398 
Flour,  wheat,  261 


536 


INDEX 


Fluorine,  102,  103 

Fluorspar,  102 

Fluosilicates,  298 

Fluosilicic  acid,  298 

Flux,  504 

Food  values,  288 

Force,  2 

Formaldehyde,  241 

Formaline,  241 

Formation,  heat  of,  281 

Formic  acid,  242 

Formic  aldehyde  fumigation,  500 

Franklin,  E.  C.,  156 

Franklinite,  396 

Fraunhofer  lines,  368 

Freezing  point,  definition  of,  44 

change  with  pressure,  44 

of  solutions,  418 
Fructose,  256 
Fuel  values,  288 
Fulminating  gold,  455 

mercury,  407 

silver,  451 
Fumaroles,  301 
Fumigating,  500 

G 

Gadolinite,  469,  471 
Gadolinium,  471 
Gahn,  304,  496 
Gahnite,  396,  462 
Galactose,  258 
Galenite  (galena),  478 
Gallium,  467 
Galvanized  iron,  398 
Garnierite,  516 
Gas,  candle  power  of,  268 

carbonum,  223 

conductivity  of,  424 

detonating,  32 

enriching  of,  267 

illuminating,  265 

laws  applied  to  solutions,  422 

marsh,  232 

natural,  232 

sylvestre,  223 

water,  224 
Gases,  absorption  of,  by  liquids,  411 

diffusion  of,  18 

law  of  combination  of,  42,  69 

spectra  of,  369 
Gasoline,  233 
Gautier,  146 
Gay-Lussac,  53,  70,  111 
Gay-Lussac,  law  of,  40,  42,  69,  152 

tower,  198 
Geber,  161 
Gelatine,  264 
Gerhardt,  77,  79 
Germanium,  472 
German  silver,  440,  516 
Germicides  in  agriculture,  439 


Gersdorffite,  516 
Glass,  383 

Bohemian,  385 

bottle,  386 

colored,  385 

composition  of,  386 

crown,  385 

cut,  385 

enamel,  386 

flint,  385 

hard,  385 

history  of,  386 

ordinary,  soda-lime,  384 

plate,  384 

potash-lime,  385 

soft,  385 

window,  384 
Glauber,  49,  150,  161 
Glauber's  salt,  46,  362 
Glazes,  466 
Glover  tower,  198 
Glucinum,  393 

compounds  of,  393 
Glucose,  256 
Gluten,  261 
Glycerine,  238 
Glycol,  238 
Glycolic  acid,  245 
Gold,  436,  452,  453 

alloys  of,  454 

analytical  tests  for,  456 

compounds  of,  454 

electroplating  with,  455 

metallurgy  of,  452 

production  of,  454 
Goldschmidt's  process,  459,  486 
Gram-molecule  defined,  133 
Granulose,  260 
Grape  sugar,  256 
Graphite,  211 
Graphitic  acid,  212 
Greenockite,  400 
Grotthus,  426,  427,  428 
Grotthus's  theory,  426 
Guajacol,  240 
Gun  cotton,  262 
Gun  metal,  440 
Gunpowder,  black,  349 

smokeless,  262 
Gypsum,  177,  380 

dead  burned,  381 


Hales,  223 

Halides  of  copper,  441 
of  lead,  481 
of  silver,  448 

Halogens,  96 

compounds  of,  with  each  other,  118 
compounds  of,  with  sulphur,  188 
general  relations  of,  to  one  another, 
119 


INDEX 


537 


Hammer  black,  510 
Hardness  of  water,  220,  376 
Hartshorn,  spirits  of,  150 
Hausmannite,  495 
Heat  of  combustion,  282 

tables  of,  286 
Heat  of  formation,  281 

tables  of,  284,  285 
Heat  of  neutralization,  282,  430 

table  of,  287 
Heat  of  solution,  49,  281 
Heavy  spar,  387 
Helium,  148,  149 
Helium  group,  discovery  of,  139 
Hematite,  502 
Henry,  411 
Heptane,  232 
Hermann,  400 
Hesiod,  503 
Hess,  277 
Hess,  law  of,  278 
Hexagonal  system,  183 
Hexane,  232 
Hippuric  acid,  244 
Hisinger,  424,  470 
Hjelm,  492 
Hjelmite,  469 
Hoffman  apparatus,  40 
Homer,  503 

Homologous  series,  233 
Hornblende,  394 
Horn  silver,  444 
Human  body,  composition  of,  9 
Hydrargyrum,  401 
Hydrates,  47 
Hydrazine,  157 
Hydrazoic  acid,  159 
Hydriodic  acid,  114,  115 
Hydrogel,  295 

Hydrobromic  acid,  107,  108,  109 
Hydrocarbons,  232 

general  behavior  of,  236 

halogen  substitution  products  of,  237 

preparation  of,  234 
Hydrocarboxylic  acids,  245 
Hydrochloric  acid,  49 

chemical  behavior  of,  51 

composition  of,  50,  52 

electrolysis  of,  51 

preparation  of,  87 

solutions  of,  50 
Hydrocyanic  acid,  230 
Hydrofluoric  acid,  104 
Hydrofluosilicic  acid,  298 
Hydrogen,  13 

absorption  of,  by  palladium,  524 

adsorption  of,  by  solids,  19 

antimonide,  324 

diffusion  of,  18 

dioxide,  91 

history  and  occurrence  of,  13 

in  the  air,  146 


occlusion  of,  by  platinum,  526 

preparation  of,  14,  15,  16,  85,  86 

properties  of,  17 

uses  of,  22 
Hydrogen  peroxide,  91 

bleaching  with,  95 

formula  of,  94 

properties  of,  92 

uses  of,  95 
Hydrogen  persulphide,  188 

selenide,  204 

silicide,  296 

sulphide,  185,  188 
Hydrolysis,  130,  131 

irreversible,  130 

reversible,  132 
Hydronitric  acid,  159 
Hydroquinone,  240 
Hydrosol,  295 
Hydroxides,  47 
Hydro  xylamine,  158 
Hygroscopicity,  46 
Hypo,  363 
Hypobromites,  109 
Hypobromous  acid,  109 
Hypochlorites,  97 
Hypochlorous  acid,  97 
Hypoiodites,  116 
Hypoiodous  acid,  116 
Hyponitrous  acid,  170 
Hypophocphoric  acid,  313 
Hypophosphorous  acid,  313 
Hyposulphite  of  soda,  202,  363 
Hypothesis  (see  Theory) 


Ice,  action  of,  on  salt,  5 

crystalline  nature  of,  45 

machines,  156 
Iceland  spar,  376 
Illuminants,  267 

Incandescent  lamp  filaments,  493 
Indicators,  130 

neutrality  toward,  131 
Indium,  468 
Ingots,  507 
Ink,  India,  215 

ordinary,  513 

sympathetic,  519 

"white,"  515 
Insulators,  424 
Inulin.  257 

Inversion  of  sugar,  256 
Invertase,  257,  259 
lodates,  117 
lodic  acid,  116         0 
Iodides,  116 
Iodine,  111 

monobromide,  118 

monochloride,  118 

oxide  of,  116 

oxy-acids  of,  116 


538 


INDEX 


Iodine  —  Continued 

pentafluoride,  118 

preparation  of,  111,  112,  113 

properties  of,  113 

tincture  of,  113 

trichloride,  118 

uses  of,  114 
lodocrol,  114 
lodoform,  114,  237 
Ion,  432 

Ionic  theory,  428 
Ions,  425,  429,  430 
Iridium,  525 
Iridosmium,  524 
Iron,  502 

action  of,  on  sulphur,  3 

alloys,  508 

analytical  tests  for,  515 

carbide,  515 

cast,  gray,  white,  505  . 

chlorides  of,  510 

dialyzed,  510 

disulphide,  512 

electrolytic,  508 

metallurgy  of,  503 

occurrence  of,  502 

ore,  502 

ore,  bog,  510 

ore,  brown,  510 

ore,  chrome,  485 

ore,  titaniferous,  299 

oxides  of,  28,  509 

passive,  509 

phosphides  of,  515 

production  of,  508 

properties  of,  508 

silicides  of,  515 

wrought,  505 
Isomers,  10,  248 
Isometric  system,  182 
Isomorphism,  law  of,  80,  185 
Isosmotic  solutions,  423 
Isotonic  solutions,  423 


Javelle  water,  99 
Jorgensen,  521 


K 


Kahtenberg,  423,  426 
Kainite,  344,  394 
Kaolin,  465     • 
Kelp,  111 
Keratine,  264 
Kermes  mineral,  330 
Kerosene,  233 
Ketones,  255 
Kieserite,  394 
Kindling  temperature,  27 
Kipp  apparatus,  219 
Kirchhoff,  366,  369 


Kjeldahl's  method,  151 
Klaproth,  205,  299,  470,  492,  493 
Kreosol,  240 
Krypton,  149 
Kunkel,  304 


Lactates,  246 
Lactic  acid,  245 

bacteria,  246 
Lactose,  258 
Ladenburg,  90 
Lsevulose,  256 
Lakes,  462 
Land  plaster,  381 
Lapis  infernis,  450 

lazuli,  467 
Lampblack,  214 
Lamy,  468 

Lana  philosophica,  398 
Lanthanum,  469 
Laplace,  277 
Latent  heat,  of  fusion,  275 

of  vaporization,  275 
Laughing  gas,  171 
Laundry  blue,  467 
Laurent,'  77 
Lavoisier,  33,  34,  35,  58,  68,  139,  161, 

223,  277 

Law  and  theory,  difference  between.  68 
Law  defined,  4 
Law  of  Henry,  411 

Hess,  278 

Lavoisier  and  Laplace,  277 

maximum  work,  279 

octaves,  338 

thermoneutrality,  430 
Lead,  478 

acetates,  482 

alloys  of,  479 

analytical  tests  for,  484 

arsenate,  482 

bromide,  481 

carbonate,  482 

chamber  process,  194 

chloride,  481 

chromate,  490 

hydroxide,  480 

iodide,  481 

nitrate,  482 

oxides,  29,  480 

oxybromide,  481 

oxychloride,  481 

peroxide,  480 

pipes,  479 

production  of,  478 

properties  of,  478 

sesquioxide,  480 

sulphate,  482 

sulphide,  482 

tetrachloride,  481 


INDEX 


539 


Lead  —  Continued 

tree,  479 

uses  of,  479 

water,  482 
Le  Bel,  247 

Le  Blanc  soda  process,  358 
Le  Chatelier,  principle  of,  45 
Lecithine,  304 

Lecoq  de  Boisbaudran,  467,  471 
Lenssen,  337 
Leucippus,  68 
Lewes,  271 
L'Hermite,  421 
Liebig,  394 
Ligroin,  233 
Lime,  377 

chloride  of  (see  Bleaching  powder) 

feldspar,  465 
Limekiln,  377 
Limestone,  376 
Limewater,  377 
Limonite,  502,  510 
Liquid  air,  147 
Litharge,  480 
Lithium,  364 

compounds  of,  364,  365 

nitride,  156 

Lithosphere,  composition  of,  8 
Litmus,  130 
Liver  of  sulphur,  352 
Lodestone,  510 
Lessen,  158 
Lubricating  oils,  233 
Lunar  caustic,  450 
Lunge,  195 
Lutecium,  470 
Luteocobaltic  chloride,  520 

M 

Magenta,  262 
Magnalium,  459 
Magnesia,  394 

alba,  395 

calcined,  394 

usta,  394 

Magnesite,  394,  395 
Magnesium,  393 

amalgam,  403 

ammonium  arsenate,  396 

ammonium  chloride,  395 

ammonium  phosphate,  314,  396 

carbonates,  395 

chloride,  395 

chromate,  490 

group,  general  remarks  about,  409 

group,  metals  of,  393 

hydroxide,  394 

metallic,  394 

nitride,  156,  394 

oxide,  394 

phosphates,  396 

pyroarsenate,  396 


pyrophosphate,  396 

silicide,  290 

sulphate,  396 

tests  for,  396 
Magnetic  iron  ore,  510 
Magnetite,  502,  510 
Malachite,  436,  443 
Malic  acid,  250 
Malleable  iron,  506 
Malonic  acid,  245 
Maltose,  258 
Manganates,  498 
Manganese,  495 

analytical  tests  for,  501 

blende,  495 

bronze,  496 

carbonate,  498 

chloride,  498 

dioxide  (peroxide),  496 

heptoxide,  497 

monoxide,  496 

nitrate,  498 

properties  of,  495 

protosesquioxide,  496 

salts  of,  497 

sesquioxide,  496 

spar,  495 

sulphate,  497 

trioxide,  497 
Manganic  acid,  498 

chloride,  498 

compounds,  495,  498 

sulphate,  498 
Manganine,  516 
Manganite,  495 
Manganites,  497 
Manganous  carbonate,  498 

chloride,  497 

compounds,  495,  496,  497,  498 

hydroxide,  496 

nitrate,  498 
Mannite,  238 
Marine  acid  air,  49 
Marl,  376,  465 
Marsh's  test,  319 
Mass  action,  132 

illustrations  of,  134 

law  of,  134 

Mass,  conservation  of,  10 
Matches,  308 
Matter,  1 
McBride,  223 
Meerschaum,  394 

Membrane,  semipermeable,  420,  423 
Mendele-eff,  338,  340,  469,  472 
Mendele"effs  view  of  Arrhenius's  theory, 

432 
Mercuric  ammonium  chloride,  408 

bromide,  406 

chloride,  405 

cyanide,  229,  406 

diammonium  chloride,  408 


540 


INDEX 


Mercuric  fulminate,  407 

iodide,  406 

nitrate,  406 

oxide,  404 

sulphate,  407 

sulphide,  black,  red,  407 
Mercurous  ammonium  chloride,  408 

ammonium  nitrate,  408 

bromide,  405 

chloride,  404 

chromate,  490 

iodide,  405 

nitrate,  406 

oxide,  404 

sulphide,  407 
Mercury,  401 

compounds  of,  403 

compounds,    physiological    properties 
of,  408 

halides  of,  404 

lamp,  402 

oxides  of,  404 

salts,   compounds  of,  with  ammonia, 
408 

tests  for,  408 
Meta-aluminates,  462 
Meta-aluminic  acid,  461 
Meta-arsenates,  321,  322 
Meta-arsenic  acid,  322 
Metaborates,  302 
Metaboric  acid,  301 
Metaformaldehyde,  248 
Metalloids,  7 
Metals  of  the  earths,  457 
Metamerism,  249 
Metantimonious  acid,  327 
Metaphosphoric  acid,  313,  315 
Metaplumbic  acid,  481 
Metastannates,  476 
Metastannic  acid,  476 
Metathesis,  9 
Methane,  232 
Methyl,  amine,  262 

butyrate,  252 

chloride,  237 

formate,  252 

hydrogen  sulphate,  252 

iodide,  252 

orange,  130 

salicylate,  252 

silicate,  299 

sulphide,  255 

Meyer,  Lothar,  51,  338,  342 
Meyer,  Victor,  189,  318 
Microcosmic  salt,  314 
Milk  of  lime,  377 
Milk  sugar,  258 
Millon's  reagent,  407 
Minium,  480 
Mirrors,  403 
Mispickel,  317 
Mitscherlich,  80,  185 


Mixture,  410 

Mohr's  salt,  513 

Moissan,  102,  103,  211,  299,  336 

Molasses,  258 

Molecular  theory,  basis  of,  72 

Molecular  weight,  determinations  of,  72 

determinations  of,  in  solutions,  417 

in  solutions,  discussion  of,  419 
Molecules,  70,  77 
Molybdates,  491 
Molybdenite,  491 
Molybdenum,  485,  491 

compounds  of,  492 

trioxide,  491 
Molybdic  acid,  492 
Monazite,  300,  469,  470 
Monocalcium  phosphate,  305 
Monoclinic  system,  184 
Monoses,  255 

Monosymmetric  system,  184 
Mordants,  462 
Morphine,  263 
Mortar,  377 
Mosaic  gold,  477 
Mosander,  469 
Moth  balls,  236 
Miiller  von  Reichenstein,  205 
Multiple  proportions,  law  of,  30,  58 
Muntz  metal,  439 
Muriatic  acid,  49 
Mycoderma  aceti,  242 


N 

Naphtha,  233 
Naphthalene,  236 
Narcotine,  263 
Nascent  state,  87 
Negatives,  499 
Neodymium,  471 
Neon,  149 
Neoytterbium,  470 
Nernst  lamp,  471 
Nessler's  reagent,  157,  406 
Neutralization,  act  of,  122 

Arrhenius's  view  of,  430 

heat  of,  282,  287 
Newlands,  338 
Newton's  metal,  331 
Nicholson,  424 
Nickel,  502,  516 

alloys  of,  516 

ammonium  sulphate,  517 

analytical  tests  for,  521 

carbonyl,  517 

coins,  516 

glance,  516 

plating,  517 

salts  of,  517 

steel,  516 
Nickelic  hydroxide,  517 

oxide,  517 


INDEX 


541 


Nickelous  chloride,  517 

cyanide,  517 

hydroxide,  516 

nitrate,  517 

oxide,  517 

sulphate,  517 

sulphide,  517 
Nicollite,  516 
Nicotine,  263 
Nilson,  469 
Niobium,  336 
Niter,  161 
Nitrates,  of  mercury,  406 

test  for,  167 
Nitric  acid,  161 

properties  of,  164 

red,  fuming,  163 
Nitric  oxide,  166 
Nitrides,  142 
Nitrobenzene,  262 
Nitrocellulose,  262 
Nitrogen,  139 

assimilation  of,  145 

compounds,     general     considerations, 
172 

compounds  of,  with  halogens,  159 

dioxide,  167 

distribution  of,  140 

in  rain  water,  146 

iodide,  160 

molecular  formula  of,  143 

oxides,  composition  of,  59 

pentoxide,  165 

preparation  of,  140,  141 

properties  of,  142 

tetroxide,  167 

tribromide,  160 

trichloride,  159,  160 

trioxide,  170 

valence  of,  157 
Nitroglycerine,  254 
Nitrosyl  chloride,  165 
Nitrosyl  sulphuric  acid,  195,  196 
Nitrous  acid,  169 
Nitrous  oxide,  171 
Nomenclature,  84,  101 
Nonane,  232 

Nordhausen  sulphuric  acid,  201 
Nucleoproteins,  264 
Nux  vomica,  263 


Occlusion,  19 

Ocean,  composition  of,  8 

Octane,  232 

Oil  gas,  267 

Oil  of  mirbane,  262 

Oleic  acid,  244 

Olefiant  gas,  235 

Olivine,  293 

Olive  oil,  252 

Onnes,  Kammerlingh,  149 


Opal,  291 

Open  hearth  process,  508 
Opium,  263 
Optical  activity,  246 
Organic  acids,  241 
Organisms  in  water,  146 
Orpiment,  317 
Orthite,  469 
Orthoclase,  294,  465 
Orthophosphoric  acid,  313 
Orthorhombic  system,  184 
Osmates,  525 
Osmic  acid,  524 
Osmiridium,  524 
Osmium,  524 

compounds  of,  524 
Osmosis,  419 

Osmotic  pressure,  419,  421 
Oxalic  acid,  244 
Oxamide,  230 
Oxidation,  21,  86 

stages  of,  28 

Oxides,  acid-forming,  120 
Oxides  of  cobalt,  518 

copper,  440 

iron,  509 

lead,  479 

manganese,  496 

nickel,  516 

nitrogen,  161 

silver,  447 

tin,  475 
Oximes,  159 

Oxy-acids  of  nitrogen,  161 
Oxygen,  24 

atomic  weight  of,  71,  74 

history,  occurrence,  24 

preparation,  86 

properties,  25 
Oxyhydrogen  blowpipe,  31 
Ozone,  88 

in  the  air,  146 

properties  of,  90 

relation  of,  to  oxygen,  89 
Ozonic  acid,  95 


Packfong,  516 
Paint,  483 
Painter's  colic,  484 
Palladium,  524 
Palladium  compounds,  524 
Palladium-hydrogen,  524 
Palmitic  acid,  244 
Paper,  261 

sizing  of,  464 
Paracelsus,  13,  408 
Paraffin,  233 
Paraffin  series,  232,  233 
Paraformaldehyde,  248 
Pararosaniline,  262 
Paris  green,  321 


542 


INDEX 


Parke's  process,  445 
Parting,  453 
Passive  state,  509,  516 
Pasteur,  251 
Pattinson's  process,  446 
Pattinson's  white  lead,  481 
Pearlash,  350 
Peligot,  493,  494 
Pentane,  232 
Pepsin,  259 
Peptones,  264 
Perchlo  rates,  100 
Perchloric  acid,  100 
Percussion  caps,  407 
Periodates,  117 
Periodic  acid,  117 
Periodic  law,  342 
Periodic  system,  337,  339 
Permanent  white,  389 
Permanganates,  498 

uses  of,  500 
Permanganic  acid,  499 
Persulphates,  202 
Persulphuric  acid,  202 
Petit,  79 
Petroleum,  232 
Petroleum  ether,  233 
Petzite,  452 
Pewter,  474 
Pfeffer,  422 
Phenacite,  393 
Phenolates,  240 
Phenolphthalein,  130 
Phenols,  240 
Phenyl  hydrazine,  158 
Phlogistic  theory,  33 
Phlogiston,  33,  53 
Phosgene,  226 
Phosphate  rock,  382 
Phosphine,  liquid,  311 

solid,  311 
Phosphines,  309 
Phosphomolybdic  acid,  492 
Phosphonium  compounds,  310 

iodide,' 310 

Phosphor  bronze,  440 
Phosphoric  acid,  120,  313 

glacial,  315 
Phosphorous  acid,  316 
Phosphorus,  304 

compounds  of,  with  sulphur,  317 

group,  general  considerations  of,  333 

in  iron,  505 

oxides  and  acids  of,  312,  316 

oxybromide,  312 

oxychloride,  312 

oxyfluoride,  312 

pentabromide,  312 

pentafluoride,  312 

pentasulphide,  317 

pentoxide,  120,  313 

preparation  of,  305 


red,  amorphous,  307 

sulphochloride,  317 

tribromide,  312 

trichloride,  311 

trifluoride,  312 

triiodide,  312 

uses  of,  307 

yellow  or  white,  306 
Phosphotungstic  acid,  493 
Photography,  448 
Physical  change,  1,  2,  5 
Physical  mixture,  3 
Pig  iron,  504 
Pineapple  oil,  252 
Pink  salt,  475 
Pintsch  gas,  267 
Pitchblende,  390,  493 
Plants,  nitrogen  supply  of,  145 
Plaster  of  Paris,  380 
Platinates,  527 
Platinic  chloride,  526 

hydroxide,  527 

oxide,  527 

sulphide,  527 
Platinous  hydroxide,  527 

oxide,  527 

sulphide,  527 
Platinum,  525 

analytical  tests  for,  527 

black,  525 

care  of,  526 

catalytic  action  of,  526 

extraction  of,  from  ores,  522 

family,  522 

sponge,  525 

use  of,  526 
Pleonast,  462 
Pliny,  479 
Plumbago,  211 
Plumbates,  481 
Plumbic  acid,  480 

oxide,  480 
Plumbites,  480 
Plumbum,  478 

candidum,  472 

nigrum,  472 
Polariscope,  249 
Pollux,  353 
Polonium,  392 
Polymerization,  248 
Polymorphism,  179 
Polysulphides,  188 
Polythionic  acids,  202 
Porcelain,  466 
Portland  cement,  378,  379 
Positives,  449 
Potash,  350 
Potash,  caustic,  346 
Potash,  feldspar,  465 
Potassium,  343 

alum,  464 

amalgam,  347 


INDEX 


543 


Potassium  —  Continued 
amide,  155 
antimonate,  328 
antimonyl  tartrate,  327 
arsenyl  tartrate,  328 
auric  cyanide,  455 
aurous  cyanide,  455 
bicarbonate,  351 
bichromate,  488 
bisulphate,  352 
bisulphite,  352 
boryl  tartrate,  328 
bromate,  348 
bromide,, 345 
carbonate,  350 
chlorate,  99,  347 
chlorite,  101 
chlorplatinate,  526 
chromate,  488 
chrome  alum,  465,  488 
citrate,  251 
cobaltic  cyanide,  520 
cobaltic  nitrite,  520 
cobaltous  cyanide,  520 
compounds  with  halogens,  345 
cyanate,  349 
cyanide,  229,  230,  349 
diuranate,  494 
ferrate,  510 
ferricyanide,  514 
ferrocyanide,  230,  513 
fluoride,  346 
fluosilicate,  351 
hydride,  344 
hydroxide,  346 
iodate,  348 
iodide,  345 
iodide,  uses  of,  114 
manganate,  498 
metantimonate,  328 
metantimonite,  327 
nitrate,  348 
nitrite,  349 
oxide,  347 

perchlorate,  100,  348 
permanganate,  499 
peroxide,  347 
persulphate,  202 
phosphates,  351 
platinic  chloride,  353 
platinous  chloride,  527 
polysulphides,  352 
pyroantimonate,  328 
pyrosulphate,  352 
ruthenate,  523 
silicate,  351 
silver  cyanide,  451 
stannite,  475 
sulphate,  351 
sulphides  of,  352 
sulphite,  352 
sulphocyanate,  231,  350 


sulphophosphate,  317 

sulphydrate,  352 

tellurate,  207 

tellurite,  207 

tests  for,  353 

thiosulphate,  352 

water  glass,  351 

/incate,  125 
Pottery,  466 
Powder  of  algaroth,  326 
Praseodymium,  471 
Precipitation,  415,  430 
Prefixes  and  endings,  102 
Preparing  salt,  476 
Preservation  of  railroad  ties,  234 
Priestley,  24,  33,  34,  49,  139,  150,  166, 

223 

Producer  gas,  224 
Propane,  232 
Properties,  chemical,  12 
Propionic  acid,  244 
Propyl  amine,  262 

chloride,  237 

iodide,  237 

tartrate,  252 

Proteins,  or  proteids,  263 
Proust,  68 
Proustite,  444 
Prussian  blue,  514 
Prussiate  of  potash,  red,  514 

yellow,  230,  514 
Prussic  acid,  230 
Ptomaines,  264 
Puddling,  505 
Purple  of  Cassius,  455 
Purpureo-cobaltic  chloride,  520 
Putrescine,  264 
Putty,  377 
Pyrargyrite,  444 
Pyridine,  263 

use  of,  in  osmosis,  423 
Pyrite,  502,  512 
Pyroantimonic  acid,  328 
Pyroarsenates,  322 
Pyroarsenic  acid,  322 
Pyroboric  acid,  301 
Pyrogallol,  or  pyrogallic  acid,  240 
Pyrolusite,  495 
Pyromorphite,  304,  478 
Pyrophosphoric  acid,  313,  315 
Pyrosulphates,  201 
Pyrosulphuric  acid,  201 


Quadratic  system,  183 
Quartation,  453 
Quartz,  291 
Quartz  glass,  291 
Quartzite,  291 
Quicklime,  377 
Quinine,  263 
Quinoline,  263 


544 


INDEX 


Radicals,  alkyl,  237 
Radio-activity,  390 
Radium,  390 
rays,  392 
salts  of,  391 

Ramsay,  139,  147,  148,  149 
Rare-earth  elements,  469 

Rayleigh,  139,  146,  147 

Reaction,  heat  of,  281 
irreversible,  135 
reversible,  135 

Realgar,  317 

Reciprocal  proportions,  law  of,  56.  59 

Red  lead,  480 

Red  ocher,  509 

Red  precipitate,  404 

Red  prussiate  of  potash,  514 

Red  zinc  ore,  396 

Reducing  agent,  21 

Reduction,  21,  87 

Refrigeration,  artificial,  156 

Regnault,  18 

Regular  system,  182 

Reich,  468 

Resin  soap,  464 

Respiration  of  plants,  31 

Respiration,  role  of  oxygen  in,  30   31 

Rhigolene,  233 

Rhodium,  523 

compounds  of,  523 

Rhodochrosite,  495 

Rhombic  system,  184 

Rhombohedral  crystals,  184 

Richter,  J.  B.,  57,  60,  68 

Richter,  468 

Rinmann's  green,  400 

Rochelle  salt,  250 

Rock  candy,  257 

Rocks,  disintegration  of,  43 

Roebuck,  194 

Roozeboom,  511 

Rosaniline,  262 

Roscoe,  223 

Roseo-cobaltic  chloride,  520 

Rose's  metal,  331 

Rosin,  464 

Rotary  power,  specific,  250 

Rouge,  509 

Rubidium,  353 

Ruby,  460 

Ruby  silver  ore,  444 

Rum,  238 

Ruthenium,  523 

compounds  of,  523 
Rutherford,  139 

S 

Saccharose,  257 
Saccharum  saturni,  482 
Safety  lamp,  miner's,  274 


Sal  ammoniac,  spirits  of,  150 

Saleratus,  150 

Salivation,  408 

Salt  cake,  358 

Salt,  definition  of,  123,  125 
formation  of,  123 
formation,  older  view  of,  124 
spirit  of,  49 

Saltpeter,  348 

Salts,  123 
acid,  126 
basic,  127 

neutral  or  normal,  126,  128 
Samarium,  471 
Samarskite,  469,  471 
Sandstones,  296 
Saponification,  253 
Sapphire,  460 
Sarcolactic  acid,  246 
Saturation,  degrees  of,  414 
Scandium,  469 
Scheele,  24,  33,  34,  53,  63,  139,  251,  304, 

318,  389,  492,  493,  496 
Scheele's  green,  321,  444 
Scheelite,  492 
Schlippe's  salt,  330 
Schonbein,  88 
Schonite,  351,  394 
Schrotter,  307 
Schweinfurt  green,  321 
Schweitzer's  reagent,  261,  439 
Sefstrom,  336 
Selenic  acid,  205 
Selenious  acid,  205 
Selenite,  380 
Selenium,  203,  209 

compounds  of,  204,  205 
Serpentine,  294,  394 
Setterberg,  353 
hot,  479 
iderite,  502 

iemens-Martin  process,  508 
ilica,  291 
ilicates,  action  of,  on  water,  296 

decomposition  of,  296 
ilicic  acids,  293,  296 
ilicic  acid,  esters  of,  299 
ilico-ethane,  297 
ilico-methane,  297 
ilicon,  290 

amorphous,  291 

carbide,  299 

chloroform,  298 

compounds  of,  with  halogens,  297 

crystalline,  291 

dioxide,  291 

tetrabromide,  298 

tetrachloride,  298 

tetrafluoride,  297 

tetraiodide,  298 
Iver,  436,  444 

analytical  tests  for,  451 


INDEX 


545 


Silver  — Continued 

bromide,  448 

carbonate,  450 

chromate,  451,  490 

chloride,  448 

cyanide,  450 

extraction  of,  from  ores,  445 

fluoride,  448 

fulminate,  451 

glance,  444 

iodide,  448 

mirrors,  451 

nitrate,  449 

nitrite,  450 

oxide,  448 

peroxide,  448 

phosphate,  315,  450 

plating,  451 

properties  of,  447 

pyro phosphate,  315 

refining  of,  446 

solvent  for,  447 

sterling,  447 

suboxide,  448 

sulphate,  450 
Smalt,  520 
Smaltite,  518 
Smithsonite,  396 
Soap,  253 
Soapstone,  394 
Soda,  358 

baking,  361 

calcined,  359 

caustic,  357 

crystallized,  359 

feldspar,  465 

washing,  359 

water,  221 
Sodium,  354 

alcoholate,  238 

alum,  464 

amalgam,  354,  403 

amide,  155 

bicarbonate,  361 

bichromate,  489 

bisulphate,  363 

bisulphite,  363 

benzoate,  244 

borate,  364 

bromide,  356 

carbonate,  358 

chloraurate,  455 

chloride,  354 

chlorplatinate,  526 

chromate,  489 

cyanide,  364 

diuranate,  494 

fluoride,  356 

formate,  242 

hydride,  354 

hydroxide,  357 

hydrosulphide,  364 

2N 


iodide,  356 

manganese  alum,  465 

metaphosphate.  314 

nitrate,  361 

nitrite,  361 

oleate,  253 

oxide,  356 

perchlorate,  100 

permanganate,  500 

peroxide,  356 

persulphate,  202 

phosphates,  361,  362 

pyroantimonate,  329 

pyrophosphate,  314 

silicate,  294,  364 

stannite,  475 

sulphate,  362 

sulphides,  364 

sulphite,  363 

thiosulphate,  202,  363 

tungstate,  492 

water  glass,  364 
Soffioni,  300 
Soil,  nitrogen  in,  150 

organisms  in,  146 
Solder,  474 

Solid  solution,  410,  415 
Solubility  curve  of,  ferric  chloride,  511 

magnesium  chloride,  395 

sodium  sulphate,  362 
Solubility  curves,  413 
Soluble  Prussian  blue,  514 
Solution,    distinguished    from    chemical 
compound,  6,  410 

heat  of,  281 

normal,  129 

saturated,  414 

standard,  129 

supersaturated,  363,  415 

unsaturated,  414 
Solutions,  410 

colloidal,  295,  416 

nature  and  kinds  of,  410 

of  liquids  in  liquids,  411 

of  solids  in  liquids,  412 

solid,  410,  415 
Solvay,  360 
Solvay  process,  360 
Sonnenschein's  reagent,  492 
Soot,  214 

Spark  spectra,  369 
Specific  heats,  table  of,  78 
Spectra  of  metals,  367 
Spectroscope,  366 
Spectrum  analysis,  365 
Spectrum,  continuous,  368 

reversed,  369 

spark,  369 
Spelter,  397 
Spiegeleisen,  496,  507 
Spinels,  462 
Spiritus  fumans  Libavii,  475 


546 


INDEX 


Stahl,  33,  53 
Stannates,  476 
Stannic  acid,  475,  476 

chloride,  475 

hydroxide,  476 

oxide,  476 

sulphide,  477 
Stannous  chloride,  474 

hydroxide,  475 

oxide,  475 

sulphide,  477 
Starch,  259 

paste,  260 

production  of,  222 

soluble,  260 
Stas,  223 
Stearic  acid,  244 
Steel,  mild,  structural,  tool,  506 
Stereochemistry,  248 
Stereoisomerism,  248 
Stibine,  324 
Stibnite,  323 
Stohmann,  289 
Stoichiometry,  laws  of,  60 
Stolzite,  492 
Stoneware,  466 
Storage  battery,  434 
Strohmeyer,  400 
Strontianite,  386 
Strontium,  386 

carbonate,  387 

chloride,  387 

compounds  of,  386 

dioxide,  387 

hydroxide,  387 

nitrate,  387 

oxide,  387 

sucrate,  258 

sulphate,  387 

sulphide,  381 
Structural  formulae,  82 
Strychnine,  263 
Sublimate,  corrosive,  405 
Sublimation,  112 
Substance,  definition  of,  2 
Succinic  acid,  245 
Sucrose,  257 
Sugar  of  lead,  482 
Sugars,  255 
Sulphates  of,  iron,  512 

mercury,  407 
Sulphides  of,  copper,  444 :  iron,  512 ; 

tin,  477 
Sulphites,  191 
Sulphocyanates,  230 
Sulphostannates,  477 
Sulphur,  176 

allotropic  forms  of,  179 

amorphous,  179 

auratum,  330 

bleaching  with,  190 

compounds  with  halogens,  188 


dichloride,  189 

dioxide,  189 

ethers,  254 

flowers  of,  178 

group,  general  considerations  of,  207 

hexafluoride,  188 

hexaiodide,  189 

in  iron,  505 

milk  of,  179 

monobromide,  189 

monochloride,  189 

monoclinic,  189 

monoiodide,  189 

occurrence  of,  176 

peroxide,  202 

plastic,  179 

precipitated,  179 

preparation  of,  178 

properties  of,  178 

rhombic,  178 

roll,  178 

sesquioxide,  192 

tetrachloride,  189 

trioxide,  120,  192,  194 

uses  of,  180 
Sulphuric  acid,  120 

action  of,  on  salt,  136 

contact  process,  193 

factory,  197 

hydrates  of,  201 

lead  chamber  process,  194 

monohydrate  of,  200 

properties  of,  199 

uses  of,  199 
Sulphuric  ether,  254 
Sulphurous  acid,  120,  191 
Sulphuryl  chloride,  203 
Superphosphate,  198,  305 
Superphosphate  of  lime,  382 
Sylvanite,  452 
Sylvite,  344 
Symbols,  chemical,  63 


Talc,  394 
Tank  waste,  359 
Tanning,  chrome,  489 
Tantalites,  336 
Tantalum,  336 
Tartar  emetic,  327 
Tartaric  acid,  250 
Taylor,  E.  R.,  227 
Tellurium,  205,  209 

compounds  of,  206,  207 
Tempering,  506 
Tennant,  525 
Terbium,  471 
Ternary  compound,  85 
Tetraboric  acid,  301 
Tetragonal  system,  183 
Thallium,  468 

compounds  of,  468 


INDEX 


547 


Th^nard,  53,  91 

Th£nard's  blue,  467 

ThSnard's  method,  483 

Theory   and    law,     difference    between, 

68 
Theory      of     electrolytic     dissociation, 

428 

Theory,  use  of,  68 
Thermite,  460 
Thermochemical  data,  283 

uses  of,  287 

Thermochemical  equations,  278,  279 
Thermochemistry,  275 

laws  of,  277 
Thionyl  chloride,  203 
Thiosulphates,  201 
Thiosulphuric  acid,  202 
Thomas-Gilchrist  process,  507 
Thomas  slag,  507 
Thomsen,  Julius,  279,  288,  432 
Thorium,  299 
Thorium  X,  392 
ThouleVs  solution,  406 
Thulium,  471 
Tin,  472 

amalgam,  403 

analytical  tests  for,  477 

Banca,  473 

block,  473 

chlorides  of,  474 

cry,  473 

pest,  473 

recovery  of,  from  tin  cans,  474 

stone,  472 

tetrabromide,  475 

tetrachloride,     conductivity    of,    424, 
425 

tetraiodide,  475 

uses  of,  474 
Tinned  iron,  474 
Titanium,  299 
Titration,  130 
Tombac,  439 
Toning  bath,  455 
Triads,  337 
Triazoic  acid,  159 
Triazoiodide,  160 
Triclinic  system,  185 
Tridymite,  291 
Trigonal  system,  184 
Trioleine,  252 
Tripalmitine,  252 
Tristearine,  252 
Trithiocarbonates,  229 
Trithiocarbonic  acid,  229 
Trypsin,  259 
Tungsten,  485,  492 

compounds  of,  492,  493 
Turmeric  paper,  130 
Turnbull's  blue,  514 
Turpeth  mineral,  407 
Type  metal,  324 


U 

Ultramarine,  466 
Uranium,  485,  493 

compounds  of,  493,  494 

glass,  494 

rays,  390 

yellow,  494 
Uranium  X,  392 
Urbain,  470 
Urea,  estimation  of,  142 


Valence,  82,  83 

Valentine,  Basil,  194,  328,  330 

Vanadium,  335 

compounds  of,  336 
Van  Helmont,  223 
Van  Marum,  88 
van't  Hoff,  247,  422 
Varech,  111 
Vaseline,  233 
Vauquelin,  393,  485 
Venetian  red,  509 
Verdigris,  438,  444 
Vermilion,  408 
Villiger,  95 
Vinegar,  242 
Vitriol,  blue,  442 

green,  512 

oil  of,  194 

white,  399 
Vivianite,  304 

W 

Wanklyn,  111 
Water,  36 

action  of,  on  rocks,  47 

as  solvent,  47 

clarifying  of,  462 

comparison  of, with  hydrogen  sulphide, 
188 

composition  of,  40 

compounds  with,  46 

contaminated,  39 

distilled,  37 

formula  of,  71 

gas,  224 

germs  in,  39 

glass,  294 

hard,  220 

hardness  of,  376 

in  animals,  36 

in  plants,  36 

metastable  condition  of,  44 

mineral,  39 

natural,  37 

of  crystallization,  46,  156 

pipes,  bursting  of,  43 

potable,  38 

properties  of,  42 

purification  of,  39 


648 


INDEX 


Water  —  Continued 

supercooled,  44 

thermal,  40 
Wavellite,  506 
Welding,  506 
Weldon  process,  497 
Welsbach,  Auer  von,  470,  471 
Welsbach  light,  272,  470 
Werner,  521 
Wheels,  grinding,  299 
Whetstones,  299 
Whisky,  238 
White 'lead,  483 

adulterants  of,  483 
White  precipitate,  408 
Whiting,  377 
Wine,  238 

spirit  of,  238 
Winkler,  472 
Wintergreen,  oil  of,  252 
Withering,  390 
Witherite,  387 
Wohler,  217 
Wolframite,  492 
Wollaston,  61,  523 
Wood  alcohol,  237 

use  of,  in  fumigating,  500 
Wood's  metal,  331 
Wood,  spirits  of,  237 
Work,  2 

Wrought  iron,  505 
Wiirtzite,  400 
Wulfenite,  478,  491 


Xenon,  149 


Yeast,  238 

Yellow  ocher,  510 

Yellow  prussiate  of  potash,  514 

Ytterbium,  470 

Yttrium,  469 

Yttrotantalite,  469 


Zinc,  396,  397 
amalgam,  403 
blende,  396 
bromide,  399 
carbonate,  398 
chloride,  398 
dust,  397 
fluoride,  399 
iodide,  399 
oxide,  398 
oxychlorides,  399 
spinel,  396 
sulphate,  399 
sulphide,  399 
tests  for,  399 
white,  398 

Zirconium,  299 

Zymase,  259 


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